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A BRIEF COURSE 
IN CHEMISTRY 


BY 

LYMAN C. NEWELL 

t{ 

PROFESSOR OF CHEMISTRY, BOSTON UNIVERSITY, BOSTON, MASS. 

AUTHOR OF 

EXPERIMENTAL CHEMISTRY, DESCRIPTIVE CHEMISTRY, 

GENERAL CHEMISTRY, PRACTICAL CHEMISTRY, 

LABORATORY EXERCISES FOR A 
BRIEF COURSE IN CHEMISTRY 


PART I MINIMUM ESSENTIALS 

PART II SUPPLEMENTARY TOPICS 




D. C. HEATH AND COMPANY 

BOSTON NEW YORK CHICAGO LONDON 
ATLANTA SAN FRANCISCO DALLAS 


I. Q 





,V^5- 


Copyright, 1929, 

By LYMAN C. NEWELL 

2 E 9 



♦ 


PRINTED IN U.S.A. 




JUN -8 1929 ©CIA 


9558 


PREFACE 


rS 

I 

0 

i 

b 

This book contains the text for a brief course in chemistry. 
The subject matter is limited to the topics for a minimum 
course in chemistry selected by a Committee of the American 
Chemical Society. 

The book consists of two parts. Part I contains the essen¬ 
tial topics only. Part II contains supplementary topics. 

Part I, about 225 pages of actual text, with nearly 150 
illustrations, includes the topics selected by the Committee 
as essential for a brief course in chemistry. These are the 
basal topics needed for any course, whatever the kind of 
school, the aim of the teacher, or the goal of the pupil. 

Part II, about 125 pages of text, with 85 illustrations (some 
full page), includes the supplementary topics selected by the 
Conunittee as suitable for special needs and additional study. 
Thes^ topics are intended solely to supplement the essential 
topics in Part I. The wide range of topics provides material 
for extending or developing a course adapted to local needs, 
to special fields, or to the future work of the pupil. 

Parts I and II are linked in two ways. First, the sections 
throughout the book are numbered consecutively and the 
topics are connected wholly or in part by numerous cross 
references. Second, at the end of most chapters in Part I 
there are references by number and title to supplementary 
sections in Part II. Hence the minimum essentials can be 
readily and appropriately extended by the supplementary 
topics. 

The Exercises and Problems at the end of the chapters 
and of many topics have been carefully chosen. Teachers 
are urged to use these adjuncts to drive home facts and 
principles not readily learned in other ways. 

iii 


IV 


PREFACE 


No course in chemistry, however brief, is effective without 
correlated laboratory work. Experiments for this course 
are in the author’s Laboratory Exercises for a Brief 
Course in Chemistry. This laboratory manual consists 
of the experiments selected by the Committee of the American 
Chemical Society, together with supplementary experiments 
for the pupil and demonstration experiments for the teacher. 
References by number and title at the end of the chapters and 
of certain topics in the textbook show the experiments in the 
laboratory manual needed for the brief course. 

The ample Index will permit teachers to follow up special 
topics or concentrate on certain fields. 

The books listed in the Appendix will enable teachers to 
add interest to their courses and to show the service of chem¬ 
istry to the home, the community, and the nation. 

Lyman C. Newell 

Boston, Mass. 

April, 1929 


CONTENTS 

PART I — MINIMUM ESSENTIALS 

CHAPTER PAGE 

I. The Field of Chemistry — What Chemists Do . 1 

II. Scope of a Brief Course — Substances — Prop¬ 
erties — Chemical Change.4 

III. Elements — Compounds — Mixtures ... 8 

IV. Oxygen.17 

V. Carbon — Carbon Dioxide — Carbon Monoxide 29 

VI. Hydrogen.49 

VII. Water.59 

VIII. Symbols and Formulas.79 

IX. Chemical Reactions — Equations — Calculations 85 
X. Valence ..94 

XI. Chlorine — Hydrogen Chloride — Hydrochloric 

Acid. .99 

XII. Nitrogen — The Air — Argon and Helium . .113 

XIII. Acids, Salts, and Bases.123 

XIV. Ions and Ionization.131 

XV. Ammonia — Ammonium Compounds .... 144 

XVI. Nitric Acid — Nitrates.152 

XVII. Sulfur — Sulfides.160 

XVIII. Sulfur Dioxide — Sulfurous and Sulfuric Acids 169 
XIX. Fuels — Illuminants — Petroleum — Flames . 179 

XX. Starch — Sugar — Acetic Acid — Alcohol . .194 

XXI. Metals in General.201 

XXII. Sodium and Its Compounds.205 

XXIII. Calcium Compounds.216 

XXIV. Iron and Steel — Iron Compounds . . . 232 

V 







VI 


CONTENTS 


PART II — SUPPLEMENTARY TOPICS 


TOPIC 

I. 

Elements — Energy — Chemical Change 


PAGE 

257 

II. 

The Metric System . 

. 

• 

259 

III. 

Oxygen . 



260 

IV. 

Carbon . 



262 

V. 

Hydrogen ..... 



264 

VI. 

Gases and Their Measurement 

. 

. 

264 

VII. 

Kinetic Molecular Theory 

. 

. 

271 

VIII. 

Vapor Pressure — Volumetric 
Water — Gay-Lussac’s Law 

Composition of 

273 

IX. 

Laws — Atomic Theory — Atomic 

Weights 

. 

278 

X. 

Equivalent Weights 

. 

. 

282 

XI. 

Gay-Lussac’s Law — Avogadro’s 
ULAR Weights — Molecules — 

Law — 
Formulas 

Molec- 

284 

XII. 

Finding Atomic Weights . 


. 

292 

XIII. 

Nitrogen Oxides 


. 

296 

XIV. 

Fuels — Flames 


. 

298 

XV. 

Soap . . . . . 



305 

XVI. 

Cellulose — Paper . 



307 

XVII. 

Food and Nutrition . 



310 


XVIIL Phosphorus — Arsenic — Antimony — Bismuth . 318 


XIX. 

Arrangement of the Elements by 
Weights and Atomic Numbers 

Atomic 

324 

XX. 

Fluorine — Bromine — Iodine 


332 

XXL 

Silicon Compounds — Glass 

, 

338 

XXII. 

Copper . 

. 

348 

XXIII. 

Magnesium — Zinc — Mercury 

, 

358 

XXIV. 

Silver — Photography — Gold 


368 

XXV. 

Aluminum — Clay and Clay Products . 

• 

375 

XXVI. 

Lead. 


385 

XXVII. 

Radium — Radioactivity .... 


390 

Appendix. 


398 

Index . 

. . . . . • 


401 

















BRIEF COURSE IN CHEMISTRY 


PART I 


MINIMUM ESSENTIALS 




BRIEF COURSE IN CHEMISTRY 


CHAPTER I 

THE FIELD OF CHEMISTRY — WHAT 
CHEMISTS DO 

1. The field of chemistry. — The field of chemistry is vast. 
Food, clothing, wood, metals, air, water — all these and many 
more are related directly or indirectly to chemistry. 

Every industry, large or small, involves chemistry in some 
way. In some industries chemicals are made — acids, 
alkalies, dyes, solvents, bleaching substances, disinfectants, 
and photographic materials. In other industries iron, steel, 
glass, cement, medicines, and fertilizers are manufactured. 
In still others metals are extracted from ores, gasoline is 
distilled from petroleum, gas and coke are made from coal, 
rubber is prepared from the sap of certain trees, wood is 
transformed into paper, and cotton is made into rayon and 
explosives. In special industries chemicals are used in many 
processes, e.g.^ in grinding, polishing, filtering, washing, set¬ 
tling, drying, bleaching. 

In the home, chemistry is used in cooking and cleaning, in 
providing water, gas, light, and heat, in protecting our houses 
from the weather. 

On the farm, chemistry is applied in the fertilizers, which 
are put into the ground to make crops grow, in the insecti¬ 
cides which are sprayed on trees and plants to kill pests. 

In the community, chemistry plays an important part in 
guarding the water supply, enforcing food and health laws, 

1 


CHAPTER II 


SCOPE OF A BRIEF COURSE — SUBSTANCES— 
PROPERTIES — CHEMICAL CHANGE 

3. What we study in a brief course in chemistry. — In a 

brief course in chemistry we limit our study to the properties 
and uses of certain substances and the changes involved in 
important chemical processes. And since chemistry is one of 
the sciences, we study a few laws and theories which sum¬ 
marize and interpret facts. In addition we study some 
practical application of chemistry illustrated by the manu¬ 
facture and utilization of substances related to our daily life. 

4. Substances. — In every-day language a substance is 
almost any kind of material. Thus, soil, flour, and water are 
substances. In chemical language, a substance is a special 
kind of material which is alike throughout. For example, 
sulfur is a substance, because each particle of pure sulfur is 
alike. Similarly, water, sugar, iron, sulfuric acid, salt, and 
aluminum are substances. Not all the materials used in 
chemistry consist of a single substance. Soil, for example, 
contains several different substances, and a chemist in de¬ 
scribing soil would say it consists of such single substances 
as sand, water, organic matter, clay, and limestone (see § 12). 

6. Properties. — Every substance has characteristics 
called properties. And if we know or learn some of these 
properties, we can recognize, identify, or describe the sub¬ 
stance. 

6. Physical properties. — Some properties are readily 
detected by observation, e.g., color, odor, taste, hardness, 
solubility, luster, and physical state (solid, liquid, or gas), 
while other properties need special apparatus for their 

4 


SCOPE 


5 


detection, e.gr., melting point, boiling point, conductivity of 
heat and electricity, crystal form, ease of liquefaction or 
solidification, and relative weight (^.e., whether lighter or 
heavier — and how much — than some standard such as air 
or water). 

These and similar properties are called physical properties. 
We often describe and distinguish substances by their 
conspicuous physical properties, especially the conspicuous 
or specific properties. For example, (1) copper is a metal, 
which is an excellent conductor of electricity; (2) water is a 
colorless, tasteless liquid, which dissolves a great many 
substances and boils and freezes at definite temperatures. 

7. Chemical properties. — In chemistry we study con¬ 
stantly another kind of properties called chemical properties. 
These are the characteristics of a substance exhibited in chem¬ 
ical changes. For example, iron rusts in moist air. That is, 
the lustrous, hard, metallic solid becomes a rusty brown, 
porous solid, unlike iron in appearance. Experiments show 
that iron rust is a different substance from iron. This trans¬ 
formation of iron into iron rust is an example of a chemical 
change. Iron also undergoes a chemical change when treated 
with other substances besides moist air, e.g., hydrochloric acid 
and bromine. If we prepared a list of the items, so to speak, 
in the chemical conduct of iron, these items would make up 
its chemical properties. 

Every substance has a set of chemical properties, which 
becomes evident when the substance undergoes a chemical 
change, ^.e., a change into one or more new substances. 
Hence chemical properties are indispensable in recognizing 
and describing substances. 

8. More about properties. — Let us illustrate the meaning 
of physical and chemical properties by studying sulfur. 

(1) Physical properties. Examination shows certain 
physical properties, (a) It is a yellow, crystalline solid 
without odor or taste. (6) It is heavier than water, because 
it sinks when placed in a vessel of water, (c) It is insoluble 
in water but soluble in another liquid called carbon disulfide. 
id) When heated, sulfur melts into a pale yellow liquid, 
which turns brown at a moderate temperature and remains 


6 


A BRIEF COURSE IN CHEMISTRY 


so until the temperature is quite high; then it becomes 
viscous like tar and finally boils, yielding a yellow smoke 
which looks much like sulfur, (e) Sulfur does not conduct 
electricity, for when introduced into the circuit of an electric 
bell, it prevents the ringing of the bell. (/) When rubbed 
briskly with a piece of silk, sulfur becomes electrified and 
attracts tiny pieces of paper. 

(2) Chemical properties, (a) If sulfur is heated to a high 
temperature, it takes fire and burns. The flame is blue, and 
an invisible, suffocating gas is detected by the odor; if 
burned long enough, all the sulfur is transformed into this 
gas. (&) If sulfur and a powdered metal, such as iron or 
copper, are mixed and heated in a test tube, the mixture 
begins to glow, and the incandescence often spreads through¬ 
out the mass even after the test tube has been removed 
from the flame. The product is neither sulfur nor iron, but 
another substance which has properties entirely different 
from those of the original sulfur and metal. 

Thus, step by step, we have 'established by observation 
and experiment the properties that serve to identify sulfur 
and to distinguish it from other substances. 

9. Chemical change. — Chemistry deals not only with 
substances and properties, but especially with the chemical 

changes which substances un¬ 
dergo or may be made to un¬ 
dergo. Let us consider a simple 
experiment. Sulfur and iron 
can be readily distinguished by 
their properties. Sulfur is a 
yellow, brittle solid, whereas 
iron is a gray, silvery solid 
which is attracted by a mag¬ 
net. If we pulverize the two 
substances, mix, and heat in a 
test tube (Fig. 3), the contents 
of the test tube glows, even after removal from the flame. 
Now if we break the tube and examine the contents, we find 
it is a hard, black, brittle solid, which has none of the prop¬ 
erties of either sulfur or iron. In other words sulfur and 



Fig. 3. — Heating a mixture of 
iron and sulfur in a test tube. 







SCOPE 


7 


iron have disappeared, so to speak, and a new substance 
has been formed. This change is an example of a chemical 
change. That is, a chemical change is one which involves 
the formation of a new substance. Some chemical changes 
are complex, but all chemical changes have this essential 
feature, viz., the formation of one or more new (f.e., different) 
substances. 

EXERCISES 

1. State characteristic properties of (o) glass, (6) gasolene, (c) water, 
{d) paper, (e) air, (/) lead, {g) soap, Qi) sugar. 

2. Define the term substance as used in chemistry. 

3. Give three examples of chemical change you have observed, tried, 
or experienced (a) in a kitchen, (6) in a garage, (c) at a fire. 

PROBLEMS 

(See The metric system, § 328) 

1. What is the abbreviation of gram, centigram, liter, meter, cubic 
centimeter, decimeter, milligram, millimeter? 

2. Express (a) 1 liter in cubic centimeters, (6) 2 1. in cc., (c) 1 meter 
in centimeters, (d) 250 cm. in dm., (e) 1 kg. in grams, (f) 250 gm. in mg., 
(g) 56.75 1. in cc., (h) 1250 cc. in 1., (t) 1 cc. in cu. m. 

3. How many cc. in (a) 1 liter, (b) 1 cu. dm., (c) 1 cu. m. ? 

4. If 1 m. of magnesium ribbon weighs 4 dg., how many mg. will 
5 cm. weigh? 

6 . Into how many pieces 5 cm. long can a glass tube 1 m. long be 
cut? 

6 . A flask holds 750 cc. Express its capacity in (a) 1., (b) cu. dm. 

SUGGESTIONS FOR LABORATORY WORK 

(References are to Newell’s Laboratory Exercises in Chemistj’y) 

Exercise 1 — Properties and Chemical Change. 

j 

SUPPLEMENTARY SECTIONS FROM PART II 

326. A good example of chemical change. 

328. The metric system (begun). 


CHAPTER III 


ELEMENTS — COMPOUNDS — MIXTURES 

10. Compounds and elements. — In chemistry we deal 
largely with compounds and elements. 

If we heat strongly a little red powder, called mercuric 
oxide, in a test tube, the powder turns dark, and minute, 
silvery drops collect on the upper part of the test tube. 
Now, if we push well down into the test tube a small glowing 
wooden splint (or a glowing joss stick), the glow increases 
and the end of the stick bursts into a flame. Clearly a 
colorless gas, which differs from ordinary air, must have been 
produced inside the tube. 

The two new substances formed from the mercuric oxide 
are the liquid, mercury, which can be seen, and the colorless 
gas, oxygen, which mingles with the air (in the test tube). 
Both substances are quite different from the original red 
solid, mercuric oxide. Therefore, we have decomposed the 
mercuric oxide into two different substances, mercury and 
oxygen — different from each other and from mercuric 
oxide. Moreover, these new substances, mercury and 
oxygen, differ from mercuric oxide in a fundamental way, 
viz., they themselves can not be decomposed by heating or 
by any other process ordinarily used by chemists — not 
easily shown in the laboratory, but nevertheless true. 

The essential difference, then, between the mercury and 
oxygen on the one hand and the mercuric oxide on the other is 
clear. Mercury and oxygen are the fundamental constitu¬ 
ents of mercuric oxide. We can decompose mercuric oxide 
into its constituents, but there the decomposition stops. 
We have reached the end, chemically speaking. These two 

8 


ELEMENTS — COMPOUNDS — MIXTURES 


9 


substances, mercury and oxygen, are examples of elements. 
Whereas mercuric oxide is an example of compounds. 

Obviously mercuric oxide is a compound of the elements 
mercury and oxygen. Hence, we may say: — 

A compound is a substance which can be decomposed into 
two or more fundamental substances called elements, or 
can be built up from elements. 

An element is a substance which we have not been able to 
decompose into simpler substances (§§ 13, 16). 

Compounds and elements may be defined in another way. 
Elements are composed of minute particles called atoms; 
the atoms of the same element are alike. Compounds are 
composed of atoms of different elements. Hence we may 
say:— ^ 

An element is a substance which consists of the same kind 
of atoms. 

A compound is a substance which consists of different 
kinds of atoms. 

In a few words, atoms are fundamental chemical units. 
Atoms unite to form molecules. If the uniting atoms are 
alike, molecules of an element are formed, e.g., oxygen gas 
consists of molecules each containing two atoms. If the 
! uniting atoms are different, molecules of a compound are 
formed, e.g., two atoms of hydrogen and one atom of oxygen 
form one molecule of water. 

11. Characteristics of compounds. — Compounds have 
several essential characteristics. 

(1) The elements in compounds are chemically united. 
That is, the elements in a compound are not merely mingled 
or just lying side by side. Only by a chemical change can 
we decompose a compound into its elements or unite elements 
into a compound. 

(2) The properties of compounds differ from the prop¬ 
erties of the elements that compose them. Thus, the blue 
solid copper sulfate is composed of three elements — the 
red metal copper, the yellow solid sulfur, and the colorless 
gas oxygen. 

(3) The most important characteristic of chemical com¬ 
pounds is their constant composition. This means that any 



10 


A BRIEF COURSE IN CHEMISTRY 


given chemical compound always consists of an unvarying 
per cent of the same elements. ^ For example, the compound 
water always contains 88.82 per cent of the element oxygen 
and 11.18 per cent of the element hydrogen. This is a general 
fact in chemistry and it is so important that it is stated as 
the law of constant composition: — 

A chemical compound has a constant composition hy weight. 

The number of compounds is very large — many thou¬ 
sands. Yet all consist of two or more elements chemically 
united — each compound having its elements in a certain, 
fixed proportion, which is characteristic of the compound. 

12. Mixtures and compounds. — Not all the materials 
used or studied in chemistry are elements or compounds. 
Certain materials contain two or more substances, which 
can be recognized by the eye (or a magnifying glass), and 
separated by sifting, dissolving, or filtering. Any material 
which consists of substances merely mixed or stuck together 
is called a mixture. Many familiar materials are mixtures, 
e.g., paint, milk, and granite. 

Mixtures have certain unmistakable characteristics by 
which they can be distinguished from compounds. 

(1) The ingredients of a mixture may vary in kind and 
proportion. Thus, soil may be largely sand, or clay, or 
organic matter, together with many other substances. 
Indeed, the kind and proportion of ingredients may vary 
widely and the mixture still have the same name. But in 
a compound the constituents are chemically united in a fixed 
proportion. 

(2) The ingredients can be readily separated. For ex¬ 
ample, the starch in flour can be separated from the other 
parts of the ground kernel by washing out the flour with 
water. But the constituents of a chemical compound can 
be separated only by some chemical process. 

(3) The properties of a mixture vary with the kind and 
proportion of the ingredients. But every specimen of a 
compound has the same properties. For example, a mixture 
of iron and sulfur may consist of any proportion of the 
ingredients, and the properties of the mixture will depend 


ELEMENTS — COMPOUNDS — MIXTURES 11 


on the proportion; whereas the compound iron sulfide is 
always alike throughout. 

We define a mixture as a substance with varying proper¬ 
ties and a variable composition. Whereas a compound is 
a substance with fixed properties and a fixed composition. 

13. More about elements. — We shall learn much about 
elements during our study of chemistry. At present we 
need know only a few characteristics. 

(1) In the chemical changes which are shown by the usual 
chemical properties, elements are not decomposed. The 
compound water, for example, can be decomposed into the 
elements hydrogen and oxygen. But neither hydrogen nor 
oxygen has been decomposed further into other substances. 
We can say, therefore, that one marked characteristic of 
elements is that they are undecomposable substances as far 
as the usual chemical changes are concerned. 

(2) It has been customary for many years to classify as 
elements those substances which have not been decomposed 
into simpler substances. It would follow that the elements 
are the primary forms of matter, so to speak, and are chiefly 
characterized by stability. This is true of many elements. 
But this stability we observe in most chemical changes does 
not necessarily mean that elements are stable under all 
conditions. It is a fact that the rare element radium is 
spontaneously decomposing into other elements, one of which 
is helium. Other elements behave similarly. So chemists 
have been led to give elements another characteristic, viz., 
a complex structure which makes some of them unstable. 
Later we shall study radium, which is one of these unstable 
elements (see §§ 623-536). For the present it will be satis¬ 
factory to regard the chemical elements as the fundamental 
substances from which compounds are formed and into which 
compounds can be finally decomposed. 

There are about ninety elements. Only a dozen are 
abundant. It is a fact that a few elements by their 
various combinations furnish most of the substances studied 
in this book. 

We must not forget that the word element is used in two 
senses. It means, first, a simple substance in the free or un- 


12 


A BRIEF COURSE IN CHEMISTRY 


combined state, e.g., iron or sulfur. It also means one of 
the special kinds of matter in a compound, e.g., iron sulfide 
contains the elements iron and sulfur in the combined state. 

14. Each element has a symbol. — Each element is 
designated by a symbol which is an abbreviation of its name. 
The symbol is usually the initial letter, or the initial and an 
appropriate letter of the name of the element. Thus, O 
is the symbol of oxygen, C of carbon, H of hydrogen, and 
Zn of zinc. If initial letters are alike, another letter is added, 
e.g.,‘C for carbon, Ca for calcium. Cl for chlorine. In some 
cases the symbol is an abbreviation of a foreign name. 
Thus, Fe is the symbol of iron {ferrum), Cu of copper {cu¬ 
prum), Hg of mercury {hydrargyrum), Ag of silver {argentum), 
Sn of tin {stannum), Pb of lead {plumbum), Na of sodium 
{natrium), and K of potassium {kalium). 

Symbols are used constantly, (1) to designate elements, 
(2) to represent 1 atom, and (3) to express the relations of 
elements and compounds in chemical changes. A complete 
list of the elements and their symbols will be found in the 
Table on the back inside cover of this book. Important 
elements and their symbols are shown in Table I. 

TABLE I — Important Elements and Their Symbols 


Element 

Symbol 

Element 

Symbol 

Element 

Symbol 

Aluminum 

A1 

Hydrogen 

H 

Phosphorus 

P 

Bromine 

Br 

Iodine 

I 

Potassium 

K 

Calcium 

Ca 

Iron 

Fe 

Silicon 

Si 

Carbon 

C 

Lead 

Pb 

Silver 

Ag 

Chlorine 

Cl 

Magnesium 

Mg 

Sodium 

Na 

Copper 

Cu 

Mercury 

Hg 

Sulfur 

S 

Fluorine 

F 

Nitrogen 

N 

Tin 

Sn 

Gold 

Au 

Oxygen 

0 

Zinc 

Zn 


15. Formulas. — Just as each element is designated by a 
symbol, so each compound is represented by a formula. A 
formula is a group of symbols — the kind and number of 
symbols of the elements in the compound. Thus, FeS is 














ELEMENTS — COMPOUNDS — MIXTURES 13 


the formula of the compound iron sulfide and H 2 O of water. 
Each formula represents 1 molecule. 

16. Interpretation of chemical change by elements and 
compounds. — In a chemical change (1) elements unite to 
form a compound, or (2) a compound decomposes into ele¬ 
ments ; in some chemical changes (3) compounds unite to 
form other compounds or (4) decompose into other com¬ 
pounds. For example, (1) the element sulfur (S) and the 
element iron (Fe) unite to form the compound iron sulfide 
(FeS), (2) the compound mercuric oxide (HgO) decomposes 
into the elements mercury (Hg) and oxygen (0), (3) the 
compounds ammonia (NH3) and hydrogen chloride (HCl) 
unite to form the compound ammonium chloride (NH4CI), 
and (4) the compound calcium carbonate (CaCOs) decom¬ 
poses into the compounds lime (CaO) and carbon dioxide 
(CO 2 ). 

17. Reactions and equations. — Chemical change is 
sometimes called chemical action. And a single chemical 
change is called a reaction or an interaction. A reaction 
may be represented by a chemical equation, thus: — 

Iron + Sulfur = Iron Sulfide 

In these simple equations, the plus sign may be read and 
and the equality sign/orm(s). Thus, briefly, iron and sulfur 
form iron sulfide. Or, more fully, the elements iron and 
sulfur when heated together undergo a chemical change 
which results in the formation of the compound iron sulfide. 

In a chemical change there is no loss or gain in the total 
weight, that is, the elements and compounds merely rearrange 
themselves in another way. They are not destroyed. 
Sometimes this general fact is stated in a brief form called 
the law of the conservation of matter, thus: — 

In a chemical change substances are not destroyed but merely 
rearranged in a different way, and hence the final weight is the 
same as the original weight. 

We can represent this fundamental fact by an equation. 
Thus, 56 parts by weight of iron always unite with 32 parts 


14 


A BRIEF COURSE IN CHEMISTRY 


by weight of sulfur and form 88 parts by weight of iron sulfide, 
and the equation is: — 

Iron + Sulfur = Iron Sulfide 
56 32 88 

This equation is read: 56 parts of iron and 32 parts of 
sulfur form (or equal) 88 parts of iron sulfide. By ‘‘ parts 
by weight ” we mean any denomination, e.g., grams, kilo¬ 
grams, pounds, tons. 

We usually use symbols and formulas in equations. Thus, 
the above equation is written : — 

Fe + S = FeS 

This is the correct equation, because Fe stands for 56 parts 
of iron and S for 32 parts of sulfur. Similarly, for the de¬ 
composition of mercuric oxide we write : — 

2 HgO = 2 Hg + O 2 
2(216) 2(200) 2(16) 

18. Weighing and measuring in chemistry. — In chem¬ 
istry we do not use the English System of weights and 
measures. A different system called the Metric System is 
used. An account of the Metric System is given in § 328. 
The units of length, weight, and volume should be learned; 
also the names, abbreviations, and relations of the denomina¬ 
tions in each division, especially weight and volume. 

19. Review and summary of Chapters II and III. — The 
special kinds of materials studied in chemistry are called 
substances. Each substance has its own set of character¬ 
istics, called properties, by which it is recognized. Physical 
properties include the characteristics readily observed under 
the usual external conditions. Whereas chemical properties 
include the characteristics revealed when substances undergo 
fundamental alteration called chemical changes — changes 
in which substances are transformed into other substances. 
Elements are simple substances in the sense that each is one 
kind of fundamental material which has not been changed 
by the usual chemical processes into any other kind of simple 
material. There are about ninety elements. Each element 


ELEMENTS — COMPOUNDS — MIXTURES 15 


is designated by an abbreviation called a symbol ; a symbol 
represents 1 atom. Compounds consist of two or more 
elements in chemical combination. There are thousands of 
compounds, and each has a constant composition, which is 
expressed by a formula ; a formula represents 1 molecule. 
In mixtures, which are often studied, the proportions of the 
ingredients are not constant, and the properties vary. In 
chemical changes there is no loss or gain in the total weight 
of the reacting substances. A chemical change, or reaction, 
is expressed by symbols and formulas in a condensed form 
called an equation. 


EXERCISES 

1 . What is an element ? A compound? In what fundamental way 
do elements and compounds differ ? Could you prepare (a) a compound 
from elements, (b) elements from a compound, (c) compounds from 
compounds, (d) elements from elements? 

2 . State and illustrate two characteristics of a chemical compound. 

3 . What are some characteristics of a mixture? 

4 . (a) How can water be distinguished from gasolene ? (6) Copper 

from iron ? (c) Glass from sand ? (d) Sugar from starch ? 

6 . Name eight important elements and give the symbol of each. 

6 . What is the formula of iron sulfide, water, mercuric oxide, carbon 
dioxide ? 

7 . What is a reaction? An equation? Give an example of each. , 

8. Interpret this equation by using the terms compound and element: 
Mercuric Oxide = Mercury + Oxygen. 

9 . State and illustrate the law of constant composition. 


PROBLEMS 

1 . Suppose exactly 3.5 gm. of iron and 2 gm. of sulfur are heated 
until the chemical change is complete. What weight of iron sulfide 
is produced? {Suggestion. See § 17 .) 

2 . Suppose 5 gm. of iron and 2 gm. of sulfur are heated until the 
chemical change is complete, {a) What weight of iron sulfide is 
produced? (6) Is any sulfur or iron left over? (c) If so, which, and 
how much ? 

3 . What per cent of iron sulfide is sulfur ? Iron? 

4 . When mercuric oxide is made, 25 parts by weight of mercury and 
2 of oxygen unite. What is the per cent of (a) mercury and (6) oxygen 
in mercuric oxide? 

6. How many grams of (o) hydrogen and (6) oxygen can be obtained 
from 150 gm. of water? {Suggestion, See § 11 .) 


16 


A BRIEF COURSE IN CHEMISTRY 


SUGGESTIONS FOR LABORATORY WORK 

(References are to Newell’s Laboratory Exercises in Chemistry) 

Exercise *2 — Mixture and Compound. 

Exercise SI — Decomposition of Mercuric Oxide. 

(See also Exercises in Weighing and Measuring, end of § 328 .) 

SUPPLEMENTARY SECTIONS FROM PART II 

326 . Distribution of the elements. 

326 . A good example of chemical change (review). 

327 . Energy and chemical change. 

328 . The metric system (concluded). 


CHAPTER IV 


OXYGEN 



20. Oxygen is an important element. — Oxygen is first 
in importance among the chemical elements. Free oxygen 
is essential to all animal 
life — without oxygen 
we would die. It is 
also necessary in burning 
fuels. 

21. Oxygen is the 
most abundant chemical 
element. —Oxygen occurs 
both free and combined. 

Free oxygen forms nearly 
21 per cent (by volume) 
of the air. Combined 
oxygen constitutes 88.82 
per cent (by weight) of 
water, about 50 per cent 
(by weight) of such com¬ 
mon rocks and minerals 
as granite, limestone, 
sandstone, clay, and sand. 

The human body con¬ 
tains about 65 per cent of 
combined oxygen, while 
vegetable matter con¬ 
tains about 40 per cent. 

22. Preparation of oxygen. — One of the first chemists 
to prepare and study oxygen was the English chemist 
Priestley (Fig. 4). He heated mercuric oxide in a glass vial 

17 


Fig. 4, — The English chemist Priestley 
(1733-1804), who was one of the first 
to prepare and study oxygen (1774). 






18 


A BRIEF COURSE IN CHEMISTRY 


and collected the gas in a vessel over water, in much the same 
way as we do to-day. 

Oxygen is more easily prepared from other oxygen com¬ 
pounds, e.g., water and potassium chlorate. 

In preparing oxygen from water (H 2 O), an electric current 
is passed through water to which sulfuric acid or sodium hy¬ 
droxide has been added. Two gases, oxygen and hydrogen, 
are liberated in separate tubes or compartments. This 
method is used to prepare oxygen (and hydrogen) on a large 
scale (Fig. 5). 



Fig. 5. — A plant for manufacturing oxygen (and hydrogen) by passing 
an electric current through water containing sodium hydroxide. 


Oxygen is most conveniently prepared in the laboratory 
by heating a mixture of potassium chlorate and manganese 
dioxide. 

The mixture is put in the test tube A (Fig. 6), and gently heated. 
The oxygen escapes through the delivery tube D into bottles previously 
filled with water and inverted over the end of the tube in the pneu¬ 
matic trough. The oxygen bubbles up into the bottle and displaces 
the water. 

Oxygen is liberated more regularly and at a lower tem¬ 
perature from a mixture of potassium chlorate and man¬ 
ganese dioxide than from potassium chlorate alone. More¬ 
over, after the experiment, all the manganese dioxide can 








OXYGEN 


19 


be recovered unchanged. A substance which affects the 
speed of a reaction without being ultimately affected itself 
is called a catalyst, and the process is called catalysis (com¬ 
pare §§ 31, 193, 224). 

Oxygen can be extracted from air. Air is changed into 
a liquid by great pressure and low temperature. Liquid 
air is a mixture of liquid oxygen P 

and liquid nitrogen. By allowing 
liquid air to evaporate at the ordi¬ 
nary pressure, the nitrogen, which is 
the more volatile of the two liquids, 
escapes more 
rapidly than the c — 

oxygen. By 
regulating the 
evaporation the 
nitrogen is sep¬ 
arated from the 
oxygen. The 
gas is com¬ 
pressed into steel cylinders (see Fig. 11). 

23. The preparation of oxygen illustrates one kind of 
chemical change. — Potassium chlorate decomposes into 
oxygen and potassium chloride. This chemical change may 
be expressed by the equation: — 

2 KCIO3 = 3 O 2 + 2 KCl 



Fig. 6.- 


- Apparatus for preparing oxygen in the 
laboratory. 


Potassium 

Chlorate 


Oxygen 


Potassium 

Chloride 


This kind of chemical change is called decomposition, 
that is, a chemical change in which a compound is broken up 
chemically into other elements, compounds, or both. 

24. Physical properties of oxygen. — Pure oxygen is 
colorless, odorless, and tasteless. (Certain impurities may 
give a slight odor and taste to the gas as prepared in the 
laboratory.) It is not very soluble in water — only 3 cc. 
dissolve in 100 cc. of water at the ordinary temperature. For- 
this reason it can be readily collected and stored over water. 

Oxygen is slightly heavier than air. One liter weighs 
1.43 grams, if measured and weighed at the temperature of 















20 


A BRIEF COURSE IN CHEMISTRY 


0° C. and 760 mm. {i.e., zero degrees as registered by a centi¬ 
grade thermometer and pressure of 760 millimeters as regis¬ 
tered by a barometer). 

25. Chemical properties of oxygen. — Oxygen unites 
with many elements and interacts with many compounds. 
This combining is often accompanied by light and heat 
(Fig. 7). ^ 

At ordinary temperatures oxygen unites slowly with 
several elements, e.g., lead and copper. We describe this 
change by saying these metals tarnish or rust, ^.e., they com¬ 
bine slowly with the oxygen of the air. 

At higher temperatures the chemical conduct of oxygen 
is conspicuous. Thus, a faintly glowing piece of charcoal 
when put into a bottle of oxygen 
bursts immediately into flame. Sulfur 
burns in air with a feeble bluish flame, 
but in oxygen the flame becomes large 
and brilliant (Fig. 7). Iron can hardly 
be made to burn in air, but if steel 
wool (matted strands of iron) is merely 
heated and thrust into a bottle of 
oxygen, the iron burns, sends off a 
shower of sparks, and often forms 
drops of molten iron (Fig. 7). 

Oxygen also reacts with many com¬ 
pounds, especially compounds like 
those in wood, paper, gasolene, fuel 
oil, and numberless compounds associated with plants and 
animals (§§231, 253). The chemical action with compounds, 
as a rule, is not direct combination with the compound itself 
as in the case of elements, but with the carbon and hydrogen 
freed by the decomposition of the compound (§ 29). 

In such experiments as just described, the oxygen itself 
does not burn. It assists the burning of the other elements. 
So if we were to describe briefly the chemical conduct of 
oxygen, we should say oxygen does not burn but assists the 
burning of other substances. The conspicuous chemical 
conduct of oxygen — not burning but assisting burning — 
is sometimes called its chief chemical property. 



Fig. 7. — Sulfur (right) 
and iron (left) burning 
in a bottle of oxygen. 










OXYGEN 


21 


26. Test for oxygen. — An experiment by which we try 
to identify an element or a compound is called testing or 
making a test. The decisive behavior of a substance under 
stated conditions is called the test for the substance. Thus, 
the test for oxygen is its chief chemical property, viz., the 
gas does not burn, but assists the burning of a glowing 
splint. 

27. The chief chemical property of oxygen illustrates 
chemical change. — A chemical change like that taking 
place between oxygen and sulfur is the combining of oxygen 
and sulfur. The product is a compound of the two elements. 
This kind of chemical change is called combination, that is, 
a chemical change in which compounds are formed by the 
chemical union of two or more substances. 

28. Oxidation and oxides. — The chemical change in 
which oxygen unites with a substance is called oxidation. 
The substance which unites with the oxygen is said to 
imdergo oxidation. Substances which furnish the oxygen 
are oxidizing agents. Free oxygen and air are oxidizing 
agents, though the oxygen for oxidation is often provided 
by certain compounds of oxygen, such as potassium chlorate 
(KCIO3), nitric acid (HNO3), lead dioxide (Pb02), and 
potassium permanganate (KMn 04 ). The compound formed 
by the union of oxygen and another element is called an 
oxide of the element. 

Oxides of different elements are distinguished by placing 
the name of the element (or a slight modification of it) 
before the word oxide, e.g., magnesium oxide, nitric oxide. 
Sometimes di-, or a similar numerical syllable, is prefixed 
to the word oxide, e.g., manganese dioxide (Mn 02 ), sulfur 
trioxide (SO3), phosphorus pentoxide (P2O5). 

29. Oxidation of compounds. — Oxidation is by no means 
li m ited to elements. Many compounds burn readily, i.e., 
combine as a whole with oxygen. Thus, carbon monoxide 
(CO) unites directly with oxygen to form carbon dioxide 
(CO 2 ). Sometimes the compound decomposes and the 
parts then unite with oxygen. Thus, the ingredients of 
gasolene are compounds of hydrogen and carbon, called 
hydrocarbons (§ 246). When gasolene burns, the hydro- 


22 


A BRIEF COURSE IN CHEMISTRY 


carbons decompose, the hydrogen burning to water (vapor) 
and the carbon to carbon dioxide. 

30. Combustion is oxidation. — During oxidation heat 
is liberated, and if the heat is intense, light is also produced. 

Different substances react with oxygen, 
le., oxidize, at different rates. If oxida¬ 
tion is slow, as in the rusting of some 
metals, the temperature does not change 
appreciably, and the heat escapes about 
as fast as it is produced (Fig. 8). If oxi¬ 
dation is rapid, heat is liberated quickly, 
the temperature rises suddenly, and the 
substance burns, often with dazzling light. 
Rapid oxidation which produces heat and 
light is called combustion (Fig. 9). In 
ordinary language combustion mean-s fire 
or burning; in chemical language it 
means rapid oxidation. 

Combustible substances are those 
which burn, e.g., wood, petroleum, coal. 
Whereas incombustible substances are 
those which do not burn, e.g., brick, glass, 
stone, cement, asbestos — all “ fire proof ” 
material. 

31. The rate of oxidation. — This de¬ 
pends on several factors. An important 
one is temperature. Most combustible 
substances do not burn at all at ordinary 
temperatures, but do so rapidly if the 
temperature is raised. 

Another factor is the form or shape of 
the substance. Shavings burn faster 
than kindling wood, and the kindling 
wood in turn faster than a log. Similarly, 
powdered coal is used in cement kilns 
where lump coal would be useless, and 
fuel oil is sprayed into a furnace instead of being dropped in. 

Still another factor is availability of oxygen. To make a 
fire burn we increase the supply of oxygen by opening the 



to show slow oxida¬ 
tion. The test tube 
A is moistened and 
iron in the form of 
filings or thread is 
put inside. The 
lower end of B 
rests in the vessel 
which is full of 
water. As the iron 
is slowly oxidized, 
water rises in B 
and partly fills the 
test tube. 













OXYGEN 


23 


draft, and to shut it down, we lessen the supply by closing 
the draft. A log once on fire burns rapidly if the wind blows 
upon it. In extracting iron from iron ore, and in one process 
of making steel, air is blown through the hot mixture. 

A useful factor is catalysis, i.e., the process of hastening a 
reaction by some substance which apparently does not enter 
into the chemical change 22, 78, 193, 224). Thus, in 
making sulfuric acid, sulfur dioxide is passed over platinum to 
hasten the oxidation of the dioxide to sulfur trioxide (§ 224). 



Fia. 9. — Rapid oxidation — a Southern California oil well on fire. 


Finally, the rate of oxidation is affected by a factor which 
might be called unusual local conditions. For example, if 
the heat liberated during slow oxidation can not escape 
readily, the temperature will rise steadily to such a point 
that the substance takes fire. Thus, oily rags carelessly 
thrown aside by painters or machinists, moist hay stored 
in a poorly ventilated barn, and soft coal kept in a pile a 
long time in the air sometimes take fire without apparent 
cause. Such fires, often unexpected and disastrous, are said 
to be due to spontaneous combustion, though they are 
simply cases of slow oxidation which becomes accelerated by 
accumulated heat. 





24 


A BRIEF COURSE IN CHEMISTRY 


32. Importance of oxidation. — Rapid oxidation is es¬ 
sential to our health, comfort, and progress. We burn 
wood, gas, oil, and coal to cook our food and heat and 
light our houses. So also we use these fuels as sources of 
energy to furnish the power in countless industries by 
which we prepare materials needed for food, shelter, and 
clothing. 

Slow oxidation is also important. For example, in breath¬ 
ing, air is drawn into our lungs. Here the oxygen forms an 
easily decomposed compound with the haemoglobin of the 
blood, which distributes the oxygen compound to the tissues 
of the body. And this oxygen slowly oxidizes the digested 
food and the worn-out tissues of the body. By this slow 
oxidation, heat is liberated which keeps the body warm 
and maintains it at the proper temperature (37° C. or 
98.6° F.) for life processes. 

The decay of organic matter is also due largely to slow 
oxidation. Water is purified by spraying it into the air where 
organic matter is slowly oxidized (§ 81). Sewage, too, is 
often sprayed into the air to hasten the decomposition of 
substances in the sewage by slowly oxidizing the carbon 
compounds to carbon dioxide and water. 

The hardening of paint is not due to drying in the popular 
sense but to slow oxidation of the linseed oil in the paint. 
Driers are added which act as catalysts in hastening the 
oxidation of the thin layer of oil, which finally makes the 
color stick to the painted surface. 

33. Combustion was first interpreted by Lavoisier. — 
The answer to the question “ What happens when a sub¬ 
stance burns? ” was delayed many years by a false theory 
called the phlogiston theory. The advocates of this theory 
believed that “ combustible substances contain a principle 
called phlogiston, and that when a substance burns, phlo¬ 
giston escapes.’^ This false theory was held until about 1775 
when the French chemist Lavoisier (Fig. 10) proved by his 
own and others’ experiments: (1) that phlogiston did not 
exist, and (2) that ordinary combustion is a process of com¬ 
bining with “ a certain substance contained in the air.” 
Soon after, he showed that this “ substance ” is identical 


OXYGEN 25 

with the gas discovered by Priestley in 1774. In 1778 
Lavoisier named the gas oxygen. 

34. Oxygen is essential to life. — Free oxygen is essential 
to all forms of animal life. As stated in § 32 in breathing, 
air is drawn into our lungs; here the oxygen of the air is 
taken up by the blood, which distributes it to all parts of 
the body, where oxidation occurs. By this slow oxidation, 
waste products are formed and heat is supplied to the body. 
Two of these waste prod¬ 
ucts are carbon dioxide 
and water vapor, which 
are exhaled from the 
lungs; water vapor is 
also given off through 
the skin. 

35. Uses of oxygen.— 

Oxygen is often admin¬ 
istered to persons who 
are too ill or weak to in¬ 
hale the ordinary volume 
of air, e.g.y in cases of 
pneumonia or shock from 
surgical operations. In 
submarine boats the oxy¬ 
gen of the air used up 
is replaced by oxygen 
released from tanks. 

Aviators and mountain 
climbers {e.g., the Mt. 

Everest parties) are 
equipped with oxygen 
tanks and breathing apparatus to supply quickly the oxygen 
needed at high elevations. 

Oxygen is used in mixtures of gases which burn with 
intense heat. Thus, acetylene (a compound of carbon and 
hydrogen (C 2 H 2 ), when burned with oxygen in a proper 
apparatus (called an acetylene torch or blowpipe) produces 
one of the hottest known flames (about 3000° C.)). 

Ordinary tools cut hard metals slowly, but the tip of the 



Fig, 10. — The French chemist Lavoi¬ 
sier (1743-1794), who overthrew the 
false theory of phlogiston, explained 
combustion correctly, and laid the 
foundations of modern chemistry. 






26 


A BRIEF COURSE IN CHEMISTRY 


oxy-acetylene flame when passed slowly across the metal 
melts (“ cuts ’’) it very quickly. Metal structures, such 
as fences, bridges, frames of buildings, “ scrapped war¬ 
ships, etc., are speedily dismantled by this flame. The fire 
department of large cities is equipped with an oxy-acetylene 
outfit for cutting a passage through steel doors of vaults or 
effecting an entrance into parts of a fireproof building. 

The oxy-acetylene fiame 
is used for welding (Fig. 
II). 

36. Saving lives by 
oxygen apparatus.— 

Oxygen is used in various 
forms of breathing ap¬ 
paratus for rescue work. 
The pulmotor, or lung- 
motor, is essentially a 
pump by which air rich 
in oxygen can be forced 
into the lungs at about 
the same rate as we 
breathe. The pulmotor 
is used to resuscitate 
persons who have been 
overcome by smoke or 
poisonous gases {e.g., 
illuminating gas) or who 
have been rendered un¬ 
conscious by drowning or 
by an electric shock. Fire 
departments and police stations are supplied with pulmotors. 

Another form of rescue apparatus can be hung from the 
shoulders like a knapsack. The man’s nose is clipped so 
he must breathe through his mouth (Fig. 12). Flexible 
tubes connect his mouth with a breathing bag (right), a 
cylinder of compressed oxygen (gas) and a regenerating can 
(left); the can contains potassium hydroxide to absorb the 
water vapor and carbon dioxide exhaled from the lungs. 

A man provided with an oxygen-breathing apparatus can 





OXYGEN 


27 


safely enter places where the air contains smoke or poisonous 
gas, and make repairs, extinguish fires, or rescue workmen 
who have been overcome. Extensive use is made of this 
kind of rescue apparatus in mine disasters. 



Fig. 12. — Man equipped with oxygen-breathing apparatus. He breathes 
through his mouth from the gas bag carried on his chest (right). The 
tank of oxygen and the regenerating can are carried on his back (left). 


EXERCISES 

1 . State three important physical properties of oxygen. 

2 . Prepare written answers to these questions: (a) Why does a 
draft of air make a fire burn well? (6) Why does a draft of air some¬ 
times extinguish a candle flame? (c) What oxides are found in the 
home? (d) How do fish obtain oxygen? 

3 . (a) Make a list of the name and symbol of each element mentioned 
in studying oxygen, (b) Make a list of the compounds mentioned in 
this chapter, (c) Learn the formula of each compound (as given). 

4 . Define (a) oxidation and (6) oxide. Name five oxides. 

PROBLEMS 

1 . How many gm. of oxygen are in a bottle holding 2.5 1. (at 0° C. 
and 760 mm.) ? 

2 . (a) How many liters (at 0° C. and 760 mm.) will 25 gm. of oxygen 
occupy? (b) How many gm. will 25 1. of oxygen weigh? 

3 . A pupil prepared five bottles of oxygen, each holding 250 cc. (at 
0° C. and 760 mm.). How many gm. of oxygen were prepared? 



28 


A BRIEF COURSE IN CHEMISTRY 


4. Water contains 88.82 per cent of oxygen. If 0.5 kg. was decom¬ 
posed, how many liters of oxygen (at 0° C. and 760 mm.) were formed? 

5. Potassium chlorate contains 39.18 per cent of oxygen. If 35 gm. 
are decomposed, (a) how many gm. of oxygen are liberated, and (b) 
how many bottles each containing 250 cc. will the gas fill? 

SUGGESTIONS FOR LABORATORY WORK 
(References are to Newell’s Laboratory Exercises in Chemistry) 

Exercise *3 — Preparation and Properties of Oxygen. 

Exercise 4 — Heating a Known Weight of a Metal in Air. 

Exercise S2 — Examples of Chemical Change in Exercise *3. 

Exercise S3 — Preparation of Oxygen (Short Method). 

Exercise S4 — Preparation of Copper Oxide. (See Exercises *6, *10, 
and SIO.) 

Exercise S5 — Slow and Rapid Oxidation — T. 

SUPPLEMENTARY SECTIONS FROM PART H 

329. Preparation of oxygen from various substances. 

330. Lavoisier’s famous experiment. 

331. Oxidation and energy. 


CHAPTER V 


CARBON — CARBON DIOXIDE — CARBON 
MONOXIDE 

37. Carbon is an important and useful element. — Next 
to oxygen, the most important element is carbon. It is 
also a useful element. Like oxygen, carbon is found both 
free and combined in nature. Free carbon is the familiar 
black solid that makes up the greatest part of hard coal, 
charcoal, coke, and graphite. And strange as the contrast 
seems, the highly prized gem called diamond is also free 
carbon — pure crystallized carbon. 

38. Carbon forms a large number of compounds. — The 
natural and manufactured compounds of carbon number 
over 200,000. Most of them are called organic compounds, 
because they are so closely related to living things. 

All living things contain carbon compounds, which by 
their chemical changes sustain life. Many common sub¬ 
stances which are products of living things are compounds 
of carbon with hydrogen and oxygen, e.g., sugar, starch, fat, 
and cotton, while others consist essentially of substances 
which are compounds of carbon, hydrogen, oxygen, and also 
nitrogen, e.g., flour, wool, and meat. Carbon united with 
hydrogen forms a large class of compounds called hydro¬ 
carbons, which are the main ingredients of illuminating gas, 
natural gas, petroleum, kerosene, gasolene, fuel oil, lubricat¬ 
ing oils, paraffin wax, and turpentine. 

The manufactured compounds of which carbon is the 
central element include dyes, drugs, medicines, photo¬ 
graphic developers, perfumes, soap, ink, and a large number 
of other substances needed for business and pleasure. 

29 


30 


A BRIEF COURSE IN CHEMISTRY 


The commonest inorganic compounds of carbon, i.e., 
those related fundamentally to non-living things, are the 
carbonates and the oxides. The carbonates are compounds 
of carbon with oxygen and a metal (such as calcium, mag¬ 
nesium or sodium). Thus, calcium carbonate (CaCOs) is 
the natural substance limestone, marble, or chalk. Sodium 
carbonate (Na 2 C 03 ) is the common substance washing 
soda; it is manufactured in large quantities. There are two 
carbon oxides — carbon dioxide (CO 2 ) and carbon monox¬ 
ide (CO). 

39. Diamond. — The purest natural form of carbon is 
diamond. As found in mines, principally in South Africa, 
diamonds are usually rough looking stones, which must be 



Cut 


Crystal 


Rough 

Fig. 13. — Diamonds. 


ground, or cut,’’ into special shapes and polished to bring 
out the luster and make them sparkle in the light (Fig. 13). 

Diamond is one of the hardest substances known, and can 
be cut ” and polished only by rubbing with diamond 
powder, which has very sharp, hard edges. 

Diamond resists the action of most chemicals. It com¬ 
bines with oxygen when the two elements are heated to¬ 
gether to a high temperature. By this experiment it can 
be shown that diamond is carbon, for when pure diamond is 
burned in oxygen, the only product is carbon dioxide. 

Diamonds are sold by a weight called a carat; 1 carat = 
200 milligrams. 

40. Graphite. — Graphite is a dark lead-colored, shiny 
solid. When rubbed or powdered, it becomes minute, 
smooth, soft, slippery scales. Hence it leaves a black mark 
on paper. This property is utilized in the lead pencil. 




CARBON 


31 


In making lead pencils, the graphite is washed free from impurities, 
ground to a fine powder, mixed with clay, and then pressed through 
perforated plates, from which the “ lead ” issues in tiny rods. These 
are dried, cut into the proper lengths, baked to remove all traces of 
moisture, and then inserted in the wooden case. Varying proportions 
of clay produce different degrees of hardness. 

Pure graphite is carbon ; but it is sometimes called “ black 
lead/’ or plumbago, because it was formerly supposed to 
contain lead. 

Unlike diamond, graphite is a good conductor of electricity, 
and for this reason it is often used to coat molds in electro¬ 
typing. It resembles diamond in its insolubility in liquids 
at the ordinary temperature. 

Graphite changes into carbon dioxide when heated in¬ 
tensely in oxygen. But it can be heated to a very high 
temperature in the air without melting or oxidizing. Be¬ 
cause of its infusibility it is sometimes used to make stove 
polish, protective paints, and the electrodes of electrical 
apparatus in which great heat is produced. It is the prin¬ 
cipal ingredient of the mixture (graphite and clay) which is 
made into graphite crucibles; these crucibles can be heated 
to a high temperature without melting or oxidizing, and cer¬ 
tain metals, e.g., crucible steel, are made in them. 

Graphite, owing to its slipperiness, is used as a lubricant, 
particularly where oil might clog, e.g., on the sliding wooden 
parts of an organ. Some varieties of artificial graphite can 
be ground into very fine particles. If ground with tannin 
(or a similar substance), and mixed with oil or water, this 
graphite remains suspended a long time. Such suspensions, 
known as oildag or aquadag, make excellent lubricants. 

Graphite is not attacked by corrosive chemicals, e.g., 
chlorine and sodium hydroxide. Hence, manufactured 
graphite, in the form of rods, plates, and slabs, is an in¬ 
dispensable article in electrochemical industries (§§ 49, 51, 
313, 411, 500). 

Native graphite occurs most abundantly in Ceylon. 
Graphite is manufactured at Niagara Falls by heating a 
special grade of hard coal in a limited supply of air in an 
electric furnace. ‘The process is electrothermal, i.e., the 


32 


A BRIEF COURSE IN CHEMISTRY 


chemical change is brought about by the intense heat pro¬ 
duced by passing an electric current through the materials 
(Fig. 14). Articles of almost any size and shape can be made 
of artificial graphite. 

41. Amorphous carbon. — This group includes coal, char¬ 
coal, lampblack, coke, and gas carbon. They are varieties 
of impure carbon. The word amorphous means without 
crystal form; it is often used to designate uncrystallized, 
or very fine, substances. 

42. Coal. — There are many varieties of coal (Fig. 15). 
Bituminous coal (soft coal) contains about 70 per cent of 



Fig. 14. — An electric furnace for making graphite. 


carbon. It burns with a smoky flame and is used as a fuel 
for steam, and to make illuminating gas and coke. Anthra¬ 
cite coal (hard coal) contains 90 per cent or more of carbon. 
It ignites with difficulty, burns with little or no flame, and 
produces considerable heat. It is used mainly for domestic 
purposes — heating and cooking — especially in eastern 
United States. Lignite, or brown coal, contains only a small 
proportion of carbon, sometimes as low as 20 per cent. It 
is used as a fuel in some localities, especially near the 
deposits. 








CARBON 


33 


Besides carbon, coal contains moisture and mineral matter; 
and soft coal, especially, contains considerable volatile 
matter. 

43. Charcoal. — This substance is obtained by heating 
wood, bones, and other organic matter in closed vessels, or 
by partially burning them in the air. More or less charcoal 
may be obtained by heating many organic substances; the 
charring, as it is called, is one test for combined carbon. 

Wood charcoal is a black, brittle solid. It burns without 
flame or much smoke, and leaves a white ash, which consists 
of mineral substances originally in the wood. It resists 



Fig, 15. — Sorting coal in a breaker at a coal mine. 


the action of moisture and many chemicals; hence fence 
posts, telegraph poles, and wooden piles are often charred 
before being put into the ground. 

Most varieties of wood charcoal are very porous and are 
good absorbers of gases. Charcoal made from fruit stones 
and cocoanut shells is dense, though porous enough for use 
in gas masks to adsorb, i.e.^ take, up or absorb, poisonous 
gases. Charcoal is sometimes used in small household 
filters to purify drinking water. Charcoal used for such a 
purpose must be frequently renewed or often heated to red¬ 
ness ; otherwise it becomes clogged and contaminated. 

Wood charcoal is made by heating wood in closed furnaces, no air 
whatever being admitted. By this method, which is called dry distilla- 



34 


A BRIEF COURSE IN CHEMISTRY 


tion, the yield of charcoal is about 30 per cent. From the condensed 
volatile matter acetic acid and methanol (or “ wood alcohol ”) are 
obtained (§§ 264, 266). 

Animal charcoal or boneblack is made by heating bones 
and animal refuse in a closed vessel. Animal charcoal made 
from bones contains only about 10 per cent of carbon, which 
is distributed throughout the porous mineral matter of the 
bone (largely calcium phosphate). Animal charcoal is used 
as a pigment, especially in making shoe-blacking. It is 
also extensively used to remove colored substances from 



Fig. 16. — A battery of by-product coke ovens at one of the Ford auto¬ 
mobile plants. 


sugar sirups; the straw colored solution is clarified, i.e., 
made colorless, by filtering through layers of bone charcoal. 

44. Coke. — This variety is made by expelling the vola¬ 
tile matter from bituminous coal, somewhat as charcoal is 
made from wood. It is left in the retorts when coal is dis¬ 
tilled in the manufacture of coal gas (§ 250). On a large 
scale it is made by heating a special grade of soft coal in huge 
closed furnaces (Fig. 16), constructed so as to save the by¬ 
products, e.g.y ammonia, tar, organic compounds, and com¬ 
bustible gases. 






CARBON 


35 


Coke is a grayish, porous solid, harder and heavier than 
charcoal. It burns with no smoke and a feeble flame. 
Its most extensive use is in the iron industry. 

46. Gas carbon. — This is a black, heavy, hard soHd, which is 
deposited inside the retorts in the manufacture of illuminating gas 
(§ 260). It is almost pure carbon. Being a good conductor of elec¬ 
tricity, it is extensively used for the manufacture of the carbon rods of 
electric hghts and for plates of electric batteries. 

46. Lampblack. — This form of carbon is prepared by burning gas, 
oil, or oily substances rich in carbon in a hmited supply of air. The 
dense smoke is finely divided carbon. 

Lampblack is one of the purest forms of amorphous carbon. It is 
used in making printer’s ink, black enamel, and certain black paints. 

47. Physical properties of carbon. — Many physical 
properties have been given in the preceding sections. 

Carbon does not melt, though it volatilizes at high tempera¬ 
tures, e.g., at the temperature of the electric furnace. It is 
insoluble in the ordinary solvents. Some molten metals 
dissolve it, especially iron, which may dissolve as much as 
1 per cent of its weight of carbon. As the solution cools 
some of the carbon separates as crystals of graphite (§ 40) 
or of diamond (§ 39). Carbon, particularly the very porous 
forms of charcoal, has the property of adsorption to a marked 
degree, that is, it takes up or adsorbs large quantities of 
gases like ammonia and poison gases, and also fine solids 
like the coloring matter in sirups. 

48. Chemical properties of carbon. — Carbon does not 
interact with acids or bases. At ordinary temperatures 
carbon is an inert element, i.e., is chemically inactive. But 
as the temperature is raised, its activity increases until at 
high temperatures it is very active. Thus, carbon heated 
with oxygen forms carbon dioxide and carbon monoxide. 

At high temperatures carbon withdraws oxygen chemi¬ 
cally from oxides. Extensive application of this property 
is made in extracting metals from ores, e.g., iron from iron 
oxide (§ 305). The chemical removal of oxygen from oxides 
(and certain other oxygen compounds) is called reduction. 
The substance bringing about the change is called a reducing 
agent. Thus, carbon is a reducing agent. 


36 


A BRIEF COURSE IN CHEMISTRY 


49. Carbon disulfide. — This substance (CS 2 ) is made by direct 
combination of carbon and sulfur. In Fig. 17, EE are the carbon 
electrodes. C is charcoal and coke is fed in at KK. Sulfur is fed in at 
SSSy and accumulates at Z. The carbon disulfide vapor escapes through 
the pipe P and is condensed in a special apparatus. 

Carbon disulfide is a highly combustible liquid (keep it away from 
flames!). The odor is disagreeable and the vapor is poisonous. It is 
used as a solvent for rubber, gums, fat, and some forms of sulfur. Large 

quantities are used to destroy weevil in 
grain and in other seeds. It is effective 
in exterminating ants, mice, moles, and 
woodchucks, if the holes are closed after 
the carbon disulfide has been poured in. 

60. Calcium carbide. — This substance 
(CaC 2 ) is made by heating a mixture of 
carbon and calcium oxide (CaO) in the 
electric furnace. With water, calcium car¬ 
bide forms acetylene (C 2 H 2 ). Acetylene, 
if burned in a special burner which admits 
much air, produces a brilhant flame, and 
is used to illuminate caves and mining 
camps. With oxygen, acetylene forms a 
mixture which burns with an intensely hot 
flame, which is used in welding and 
“ cutting ” metals (§ 36). 

61. Silicon carbide. — This substance 
(SiC), better known as carborundum, is 
made by heating silica (silicon dioxide. 

Fig. 17 . — Electric furnace Si02) and coke in the electric furnace. It 
for making carbon disulfide is an extremely hard substance, nearly as 
from carbon and sulfur. hard as diamond. Its hardness has led to 
extensive use as an abrasive, e.g., grinding 
wheels, cutting stones, polishing cloth, and special shapes used particu¬ 
larly in the automobile industry (Fig. 18). 

62. Carbon is an allotropic element. — Carbon exists 
in three different modifications — diamond, graphite, and 
amorphous carbon (typified by charcoal prepared from 
sugar). All pure forms of these different substances are 
carbon. They can be changed into one another. Indeed, 
graphite is manufactured from coal (§40), diamonds have 
been made from pure charcoal, and carbon becomes graphite 
in cast iron (§ 306). Each modification burns in oxygen 
and yields only carbon dioxide. Furthermore, a given 
weight of each, say 12 gm., yields the same weight of carbon 

































CARBON 


37 


dioxide (44 gm.). Why are they so different? Recent 
experiments show that diamond and graphite have a dif¬ 
ferent structure, or in other words, they are built from the 
same material but built in a different way. 

63. Carbon and energy. — When carbon and certain 
carbon compounds (e.gr., carbon monoxide and hydro¬ 
carbons (§ 246)) burn, considerable heat is liberated. This 
means that the chemical energy stored in the carbon and 
its compounds is transformed in part into heat energy. 
Since many forms of carbon and many kinds of carbon com- 



Fig. 18. — Articles made of carborundum. 


pounds can be readily obtained, they are extensively used as 
fuels. (See Chapter XIX.) 

Solid fuels are essentially free carbon, though all, especially 
soft coal, contain other substances, e.g., moisture, volatile 
substances, and mineral matter (the last is left as ashes after 
the coal is burned). Liquid fuels are mixtures of hydro¬ 
carbons, i.e., compounds of carbon and hydrogen. Gaseous 
fuels contain hydrocarbons, carbon monoxide, and hydrogen. 

Fuels burn when heated to the proper temperature in a 
current of air. Chemically this means that the free hydrogen, 
the carbon and hydrogen from the decomposed compound, 
and the carbon monoxide combine with oxygen. Physically 






38 


A BRIEF COURSE IN CHEMISTRY 


this means that heat is liberated. Indeed, the essential 
characteristic of a good fuel is its heat-producing capacity. 

CARBON DIOXIDE 

64. Formation of carbon dioxide. — Carbon dioxide is 
formed when carbon burns in oxygen and also when a candle 
and other combustible substances burn in air. The process 
called burning or combustion consists usually in the union 
of carbon (and also, of course, hydrogen, if present) with 
oxygen (§ 33). This means that carbon dioxide is being 
formed constantly by the burning of such common fuels as 
wood, paper, coke, coal, charcoal, oil, and gas. In fact, 
carbon dioxide is always one of the products of combustion, 
as they are often called, yielded by burning any substance 
which contains carbon, e.g., sugar, starch, wax, meat, milk, 
camphor, alcohol, oil, dyes, fat, and drugs. 

The equation for the combustion of carbon is: — 

C + O 2 = CO 2 

Carbon Oxygen Carbon Dioxide 

The presence of carbon dioxide in the products of combustion 
can be shown by bubbling the products, e.g., smoke, through 
calcium hydroxide solution, or merely by shaking the prod¬ 
ucts with the solution. The calcium hydroxide becomes 
milky owing to the formation of insoluble calcium carbonate. 
This chemical change is a test for carbon dioxide (from any 
source). We express the chemical change thus : — 

CO 2 + Ca(OH )2 = CaC 03 + H 2 O 

Carbon Dioxide Calcium Hydroxide Calcium Carbonate Water 

■ The two main products of the final chemical change of 
our digested food are carbon dioxide and water. The car¬ 
bon dioxide and some water vapor are exhaled from the 
lungs. The presence of carbon dioxide in exhaled breath 
may be readily shown by blowing gently through a glass 
tube into a bottle containing calcium hydroxide (Fig. 19). 

Carbon dioxide is also formed by other chemical changes, 
such as the decay of animal and vegetable matter and the 
fermentation of organic substances like sugar. The latter 


CARBON DIOXIDE 


39 


process is illustrated by the liberation of carbon dioxide in 
making bread. 

55. Preparation of carbon dioxide. — We can prepare 
carbon dioxide by burning carbon, or a combustible com¬ 
pound of carbon, in oxygen or in an ample 
supply of air. Industrially the gas is 
manufactured by passing air through hot 
coke, purifying the gaseous product, and 
finally extracting the carbon dioxide by a 
special cooling process. The gas is used 
in preparing carbonated beverages. 

In the laboratory carbon dioxide is most con¬ 
veniently prepared by the interaction of an acid 
and a carbonate. The apparatus shown in Fig. 20 
can be used for this experiment. Dilute hydro¬ 
chloric acid and calcium carbonate (in the form 
of marble chips) are usually used. When the acid 
is introduced through B upon the calcium carbon¬ 
ate in A, the gas is rapidly liberated, bubbles up 
into bottles (previously filled with water), and dis¬ 
places the water. 

The equation for the chemical 



Fig. 19. — Blowing 
through calcium 
hydroxide to show 
the presence of 
carbon dioxide in 
the breath. 


change in the preparation of carbon dioxide is: 


+ 2HC1 

Hydrochloric 

Acid 



= CO 2 + CaCb -f H 2 O 

Carbon Calcium Water 
Dioxide Chloride 

66. Physical prop¬ 
erties of carbon 
dioxide. — Carbon 
dioxide is a color¬ 
less, odorless gas. It 
is about 1.5 times 
heavier than air. A 
liter of the pure gas 
weighs 1.98 grams (at 
0° C. and 760 mm.). 
It dissolves in water. At ordinary temperature and pres¬ 
sure, water dissolves about its own volume of carbon diox-, 
ide. Under increased pressure the solubility increases. 

The solution of carbon dioxide called soda water is manu¬ 
factured by dissolving carbon dioxide in water. A pressure 


Fig. 20. — Apparatus for preparing carbon di¬ 
oxide from calcium carbonate and an acid. 






















40 


A BRIEF COURSE IN CHEMISTRY 


of 3 to 4 atmospheres {i.e.j 3 to 4 times 760 mm.) is used. 
When soda water is drawn from a soda fountain or siphon 
(Fig. 21), the water bubbles and forms a froth, and some gas 
escapes owing to the diminished pressure. Many beverages, 
such as ginger ale, are “ carbonated,’’ ^.e., 
they are manufactured by forcing carbon 
dioxide into the prepared liquid; the bottle 
vJ is closed tightly with a cap. When the 
bottle is opened, some of the gas escapes 
(Fig. 22). 

67. Liquid and solid carbon dioxide. — Carbon 
dioxide can be readily liquefied and solidified. If 
enough pressure (about 50 atmospheres, i.e., 50 X 
760 mm.) is applied at ordinary temperatures, the 
gas becomes a hquid. Liquid carbon dioxide is 
stored and sold in strong metal cylinders. If a 
cylinder is properly opened, part of the escaping 
Fig. 21.—A siphon liquid evaporates quickly and removes so much 
of soda water. heat that the remainder soon becomes white, snow- 
hke, solid carbon dioxide. Solid carbon dioxide is 
an article of commerce, and in some places is delivered in large cakes 
like ice. It is called “ dry ice.’' It is 
also used as a refrigerant in packing and 
shipping ice cream, butter, eggs, fish, 
fruits, vegetables, and other perishable 
goods (Fig. 23). 

58. Chemical properties of car¬ 
bon dioxide. — Carbon dioxide, 
unlike oxygen, does not assist com¬ 
bustion. Nor does it burn. This 
negative behavior, so to speak, 
is sometimes called inertness or 
stability. Carbon dioxide is an 
inert gas. 

However, carbon dioxide does Fig. 22. — Carbon dioxide es- 

react with some substances. Thus, raping from a bottle of a car- 
., , . ... i.e bonated beverage. 

it combines with water to form a 




compound called carbonic acid (H2CO3) ; but this compound 
is not stable, i.e., it decomposes readily and re-forms water 
and carbon dioxide. 























CARBON DIOXIDE 


41 


We have called attention several times to the reaction 
in which carbon dioxide and calcium hydroxide form calcium 
carbonate and water. This chemical change, we have also 
said, serves as a test for carbon dioxide. A similar reaction 
takes place between carbon dioxide and sodium hydroxide. 
The product, in this case, however, is sodium carbonate 
(Na 2 C 03 ), which remains 
dissolved in the water. 

The solution feels like 
wet soap; in fact, sodium 
carbonate is sometimes 
called washing soda and 
is used in large quanti¬ 
ties as a cleansing agent. 

Sodium bicarbonate 
(NaHCOs) is closely 
related to sodium car¬ 
bonate. Sodium bicar¬ 
bonate is cooking soda. 

Alone, or as an ingredient 
of baking powder, it is 
widely used in cooking 
because it gives off car¬ 
bon dioxide which puffs 
up the dough. 

At high temperatures 
carbon dioxide reacts 
with carbon to form mon¬ 
oxide, thus: — 

CO 2 + C = 2 CO 

Carbon Dioxide Carbon Carbon Monoxide 

This chemical change takes place in a coal fire (§ 64). 

59. Relation of carbon dioxide to life. — Carbon dioxide 
is not poisonous, though the presence of a small quantity 
in the air of a room is objectionable. As already stated, 
the carbon dioxide that is exhaled from our lungs is one of 
the products formed by the oxidation of the tissues of the 
body, new tissue itself being formed from the food (§ 34). 
The carbon needed for the rebuilding of tissue is supplied 










42 


A BRIEF COURSE IN CHEMISTRY 


by fat, meat, sugar, starch, and other foods we eat. Carbon 
dioxide is a waste product of animal life. 

On the other hand, carbon dioxide is an essential food 
of plants. Through their leaves, especially, they absorb 
carbon dioxide from the atmosphere, and through their roots 
they take up water from the soil. These two compounds 
are transformed by a series of complex changes into oxygen 
and organic compounds, e.g., starch, sugar, and cellulose. 
The sunlight and the green coloring matter (called chloro¬ 
phyll) aid the plant in the formation of these compounds. 

The relation of carbon dioxide to life is clear. (1) Plants 
absorb carbon dioxide and transform it mainly into starch, 
whereas (2) animals eat starch, or similar food, assimilate it. 

Carbon dioxide Oxygen 



Fig. 24. — Cycle of carbon (A) and oxygen (R). 

and oxidize the carbon to carbon dioxide, which is exhaled 
into the atmosphere ready for the plants again, and so on. 

The significant relation of carbon dioxide and oxygen to 
plants and animals, which is often spoken of as the cycle of 
carbon and oxygen, is shown in Fig. 24. 

60. Carbon dioxide and fire extinguishers. — Carbon 
dioxide does not burn, but extinguishes burning substances. 
A saturated solution, instead of the gas itself, is frequently 
used to put out small fires. The solution is prepared, as 
needed, in portable fire extinguishers and in chemical engines 
by the interaction of sulfuric acid and sodium bicarbonate. 
The ordinary fire extinguisher contains a solution of sodium 
bicarbonate and a loosely stoppered bottle of concentrated 
sulfuric acid. When the extinguisher is to be used, the tank 
is inverted, the glass (or porcelain) stopper of the acid bottle 


CARBON MONOXIDE 


43 


falls out, the acid mixes with the sodium bicarbonate solution, 
and the pressure of the generated gas forces the saturated 
solution of carbon dioxide out of the nozzle of the extinguisher 
(Fig. 25). Some of the carbon dioxide itself escapes. The 
water solution of carbon dioxide (together with the gas), if 
directed upon the base of the fire, forces the oxygen of the 



Fig. 25. — Portable fire extinguisher. Right — cut vertically. Middle 
— partly open, showing stoppered acid bottle in original position. Left — 
inverted, showing gas bubbles and escaping water. 

air away from the fire and reduces, or entirely prevents, 
combustion. 


CARBON MONOXIDE 

61. Carbon monoxide differs from carbon dioxide. — 

Carbon monoxide (CO), like carbon dioxide, is a compound 
of carbon and oxygen, but the two compounds differ in prop¬ 
erties and composition. 

Carbon monoxide resembles carbon dioxide in being a 
gas without color, odor, or taste. But in other properties 
the two gases differ. Thus, carbon monoxide is only slightly 
soluble in water, and does not form a compound with water. 
It is lighter than carbon dioxide; a liter weighs 1.25 gm. 
























44 A BRIEF COURSE IN CHEMISTRY 

62. Carbon monoxide is poisonous. — Carbon monoxide 
is an active poison. Moreover it is very dangerous because 
the lack of odor prevents its detection. A small fraction of 
one per cent of this gas in the air produces a stupefying effect ; 
1 part in 2000 parts of air soon causes unconsciousness, and 
1 part in 750 to 800 parts of air will cause death in about half 
an hour. Many deaths have been caused by breathing air 
containing this gas. 

Carbon monoxide is an active poison because it forms a 
stable compound with the red coloring matter (haemoglobin) 


Fig. 26. — Miners testing the air in a mine for carbon monoxide. 

of the blood and thereby reduces the amount of oxygen car¬ 
ried by the blood to all parts of the body (§ 34). Hence 
persons who have been poisoned by carbon monoxide can 
not usually be revived by air, as in the case of suffocation 
by carbon dioxide and other gases. The pulmotor is some¬ 
times used to revive persons who have been overcome by 
gases containing carbon monoxide (§ 36). 

Carbon monoxide is an ingredient of ordinary illuminating 
gas, and also of the gas given off by a charcoal or a coal fire. 




CARBON MONOXIDE 


45 


Care should always be taken to prevent the escape of illu¬ 
minating gas and coal gas into rooms occupied by human 
beings. It is one of the gases produced in mine fires and 
explosions, and miners wear gas masks in rescue work after 
an accident (Figs. 26, 27). The exhaust gases from an 
automobile engine contain 
carbon monoxide, and special 
precaution should be taken 
not to run the engine in a 
garage or in a poorly venti¬ 
lated place. 

63. Chemical properties of 
carbon monoxide. — Carbon 
monoxide, unlike the dioxide, 
burns in air or oxygen. The 
flame is blue. The product is 
carbon dioxide. The equa¬ 
tion for this chemical change 
is: — 

2 CO + O 2 = 2 CO 2 

Carbon Oxygen Carbon 

Monoxide Dioxide 

The flickering bluish flame 
often seen on the top of a coal 
fire is caused by the burning 
carbon monoxide. Fig. 27. — Miner wearing a gas 

Not only does carbon mon- I™®*' examining his safety 
oxide unite readily with oxy¬ 
gen, but at a high temperature it withdraws oxygen from 
oxides. We have already seen that carbon itself acts in 
the same way (§ 48). We also learned that this chemical 
removal of oxygen is called reduction. For example, in the 
manufacture of iron from iron ores, the ore, which is an 
oxide (Fe 203 ), is reduced by carbon monoxide in a blast 
furnace, the gas for this purpose being produced by the 
incomplete combustion of coke. 

64. The two carbon oxides are formed in a coal fire. — 
The oxygen of the air entering at the bottom of a coal fire 
through the damper at A (Fig. 28) combines with the hot 






46 


A' BRIEF COURSE IN CHEMISTRY 


carbon of the coal and forms carbon dioxide — the first 
change. But the carbon dioxide in passing up through the 
upper layer of hot coal is reduced by the carbon to carbon 
monoxide — the second change. The carbon monoxide 
rises through the top of the fire into the air, where some of 
it escapes up the chimney through the damper at (7, though 
much, sometimes all, of it burns with a blue flame and pro¬ 
duces carbon dioxide — the third change. Therefore we see 
that the two oxides are closely related chemically and change 
into one another readily, especially in a coal fire. 

65. Formation of carbon monoxide. — Carbon monoxide 
is always formed when carbon and certain carbon compounds 



2 CO + O2 = 2CO2 

Carbon Dioxide 

CO 2 + C = 2 CO 

Carbon Monoxide 

c + O2 = CO2 

Carbon Dioxide 


Fig. 28. — The two oxides of carbon are formed in a coal fire. 


(e.gf., gasolene) burn in a limited supply of air. It is also 
formed when steam is passed through a hot fire of hard coal 
or coke. The gaseous product which is called water gas is 
essentially a mixture of carbon monoxide (40 to 50 per cent) 
and hydrogen (45 to 50 per cent). It burns with a hot flame 
and is used as a source of heat in industrial plants. When 
this mixture is enriched by vapor from petroleum oil so that 
it burns with a yellow flame it is used, alone or with other 
gases, as illuminating gas. Recall that the carbon mon¬ 
oxide makes such a gas poisonous. 

If air is forced through a hot coke fire, or better still if 
poor-grade coal {e.g., one leaving considerable ash) is burned 
with a limited supply of air in a special apparatus called a 
gas-producer (Fig. 29), a mixture of carbon monoxide (about 
40 per cent) and nitrogen (about 60 per cent) is formed. 
It is usually called producer gas. If both air and steam are 


















CARBON MONOXIDE 


47 


passed through hot carbon, the gaseous product contains 
hydrogen besides the carbon monoxide and nitrogen; it is 
called modified producer gas. Producer gas is easily made, 
burns uniformly and regularly 
with liberation of considerable 
heat, leaves no ashes, and is 
therefore used extensively as 
a fuel in industrial processes, 
e.g., in making lime (§ 293) 
and open-hearth steel (§ 311). 

66. Preparation of carbon mon¬ 
oxide. — Carbon monoxide is pre¬ 
pared in the laboratory by decom¬ 
posing oxalic acid with hot sulfuric 
acid, thus: — 

C 2 H 2 O 4 = CO + CO 2 + H 2 O 

Oxalic Carbon Carbon Water 

Acid Monoxide Dioxide 

The carbon dioxide is removed by 
passing the gases through a solution 
of sodium hydroxide, and the carbon 
monoxide is collected in bottles in 
the usual way. 


EXERCISES 

1. State the uses of (a) diamond, 

(6) graphite, (c) coke, {d) charcoal. 

2. How is carbon dioxide related 
to plant and animal Hfe ? 

3. In what ways can carbon di¬ 
oxide be prepared? Give equations for the reactions. 

4. What is the relation of carbon dioxide to {a) respiration, (6) fer¬ 
mentation of sugar, (c) decay, {d) making lime, (e) combustion ? 

6 . Carbon and carbon monoxide are reducing agents. Explain, and 
illustrate by equations. 

6 . Describe fully the action of carbon dioxide on calcium hydroxide. 
State the reaction by an equation. 

7. What is the test for (a) carbon, (b) carbon monoxide, (c) carbon 
dioxide, (d) a carbonate? 

8. State the equation for (a) the oxidation of carbon to carbon mon¬ 
oxide and (b) the reduction of carbon dioxide to carbon rnoooxide, 

9. Define and illustrate reductioo, 



Fig. 29. — Apparatus for manufac¬ 
turing producer gas. Coal enters 
through A, B, and air through D 
into C. Producer gas escapes 
through F, and ashes are removed 
at E. 



















48 


A BRIEF COURSE IN CHEMISTRY 


PROBLEMS 

1. What is the weight of 10 liters of carbon dioxide? 

2. A pupil prepared 5 bottles of carbon dioxide each holding 250 cc. 
How many gm. of carbon dioxide were prepared? 

3. Marble, if treated with acid, yields 44 per cent of its weight as 
carbon dioxide. If 0.5 kg. of marble was used, how many (a) gm. and 
(6) 1. of carbon dioxide were formed? 

4. What weight of marble, 90 per cent pure, is needed to prepare 150 
1 . of carbon dioxide? 

SUGGESTIONS FOR LABORATORY WORK 

(References are to NewelFs Laboratory Exercises in Chemistry) 

Exercise *5 — Properties of Charcoal. 

Exercise *6 — Reduction of Copper Oxide by Carbon. 

Exercise *7 — Preparation and Properties of Carbon Dioxide. 
Exercise S6 — Combustion and Carbon Dioxide. 

Exercise S7 — A Fire Extinguisher and Carbon Dioxide — T. 
Exercise 8 — Preparation and Properties of Carbon Monoxide — T. 
Exercise S39 — Carbonic Acid — T. 

Exercise *48 — Distillation of Soft Coal. 

Exercise *49 — Distillation of Wood — T. 

SUPPLEMENTARY SECTIONS FROM PART H 

332. How coal was formed. 

333. Carbon and energy. 

334. Carbon dioxide and plants. 

(Selections may be made from 383-392.) 


CHAPTER VI 


HYDROGEN 

67. Occurrence of hydrogen. — Hydrogen, like oxygen, 
is a gaseous element. It does not occur free to any great 
extent in nature; natural gas — a mixture of combustible 
gases which issues from the earth in certain localities — 
contains about 2 per cent. Contrary to popular belief, the 
air contains only traces of hydrogen. Vast quantities of 
free hydrogen exist in the gases that surround the sun. 
During an eclipse of the sun gigantic streamers of burning 
hydrogen may be seen shooting out thousands of miles from 
the sun’s disk into space. Mixtures of gases, used for heat¬ 
ing, and sometimes lighting, e.g., water gas and producer 
gas (§ 65), contain from 35 to 50 per cent of hydrogen. 

Combined hydrogen is a constituent of many important 
compounds. Water is 11.18 per cent hydrogen. The human 
body is about 10 per cent hydrogen. Hydrogen is a con¬ 
stituent of all acids and bases — important classes of com¬ 
pounds to be studied later. (See Chapter XIH.) 

All plants and animals contain compounds of hydrogen 
with carbon and oxygen, and in some cases with nitrogen 
also. Hydrogen is likewise a constituent of many com¬ 
pounds obtained from animal and plant products, e.gr., fat, 
meat, sugar, starch. 

Compounds of hydrogen and carbon form a large and 
important class of compounds called hydrocarbons, which 
are ingredients of petroleum (and its products, such as kero¬ 
sene, gasolene, paraffin, and lubricating oils), illuminating 
gas, and water gas. 

68. Preparation of hydrogen from acids. — Hydrogen is 
readily prepared by the interaction of certain metals and 

49 


50 


A BRIEF COURSE IN CHEMISTRY 


compounds of hydrogen called acids. The metals are zinc, 
iron, or magnesium. Dilute water solutions of the acids 
are used, e.gr., sulfuric acid (H2SO4) or hydrochloric acid 
(HCl). The hydrogen comes from the acid. The metal 
combines with the rest of the acid to form a compound called 
a salt, which usually remains dissolved in the liquid. Thus, 


Acid Metal Hydrogen Salt 
H2SO4 + Zn = H2 + ZnS 04 

Sulfuric Acid Zinc Hydrogen Zinc Sulfate 


In the laboratory hydrogen is usually prepared in a small generator, 
and collected over water in a pneumatic trough (Fig. 30). Zinc is put 
in the bottle A and acid is introduced through the dropping tube B 
by pressing the clamp. The hydrogen passes out through the dehvery 
tube D into the pneumatic trough, bubbles up into the bottles, and dis¬ 
places the water. No flame should be near during the preparation of 
hydrogen, because mixtures of hydrogen and air explode 
when ignited! 

69. Preparation of hydrogen from water. — 

Water is a compound of hydrogen and oxygen, 

and hydrogen can be 
readily prepared from 
water. 

One method has al¬ 
ready been described, 
viz., the electrolysis 
of water containing a 
little acid or alkali 




Fig. 30.—Apparatus for preparing hydrogen /s OQ) 

in the laboratory. a i • 

Another method is 

the interaction of cold water and certain metals. Potassium 
and sodium interact rapidly and calcium interacts slowly. 
Potassium interacts so rapidly that the heat ignites the lib¬ 
erated hydrogen, and the hydrogen flame is colored pale 
violet by the burning vapor of the metal (Fig. 31). But 
if a small piece of sodium is dropped upon cold water, the 
sodium melts into a shining globule, which spins about . 
rapidly on the water with a hissing sound and finally disap¬ 
pears with a slight explosion. The heat produced is not 
sufficient to ignite the hydrogen, unless the water is hot 

















HYDROGEN 


51 


or the sodium is floated on a piece of filter paper. Then 
the hydrogen flame is colored yellow by burning sodium. 
If sodium is wrapped in a piece of fine wire gauze, or of 
tea lead pierced with small holes, and 
then dropped into a dish of water, 
the hydrogen gas can be collected in 
an inverted test tube full of water 
(Fig. 32). The other product of the 
reaction is sodium hydroxide (NaOH). 
It is soluble in water, and hence is 
not seen. But if a piece of red 

litmus paper is floated on the water, 

Fig/31. — The interaction ,, • i i ii i 

of water and potassium the paper IS turned blue by the 

produces enough heat to sodium hydroxide. The equation for 

drogL^^^ liberated hy- interaction of Water and sodium 



2 H 2 O + 2 Na = H 2 + 2 NaOH 

Water Sodium Hydrogen Sodium Hydroxide 

Other metals, e.g., magnesium, zinc, and iron, liberate 
hydrogen from water only at high temperatures. Thus, 
hydrogen is liberated when steam — the 
gaseous form of water — is passed over 
heated iron. This famous experiment was 
first performed by the French chemist 
Lavoisier in 1783, while he was studying 
the composition of water. He passed 
steam through a red-hot gun barrel con¬ 
taining bits of iron. The oxygen of the 
steam combined with the iron, and the 
hydrogen escaped from the tube. Since 
Lavoisier was then studying the composi¬ 
tion of water, he named the gas hydrogen, 
which means “ water former.’^ 

70. Preparation of hydrogen from 
bases. — Hydrogen can be prepared by 
boiling a solution of certain bases with some elements. 
Bases are compounds of hydrogen, oxygen, and a metal. 
They are sometimes called hydroxides. Hydrogen is liber- 



Fig. 32. — Collecting 
hydrogen formed 
by the interaction 
of water and so¬ 
dium. 













52 


A BRIEF COURSE IN CHEMISTRY 


ated if sodium hydroxide (NaOH) is boiled with aluminum. 
The equation is as follows: — 

6 NaOH + 2A1 = 3 H 2 + 2 NagAlOs 

Sodium Hydroxide Aliiminiun Hydrogen Sodium Aluminate 

71. Commercial hydrogen. — On a large scale hydrogen 
is prepared by the electrolysis of water containing an acid 
or base (§§ 22, 97), by the interaction of iron and sulfuric 
or acetic acid, and by the reaction between steam and carbon 
in the manufacture of water gas (§ 387). Considerable is 
obtained as a by-product in the manufacture of chlorine 
and sodium hydroxide by the electrolysis of sodium chloride 
(§§ 121, 285). 

72. The preparation of hydrogen illustrates displacement. 

— The preparation of hydrogen by the interaction of a metal 
with an acid, water, or a base illustrates displacement, or, 
as it is sometimes called, substitution or replacement. In 
the case of zinc and sulfuric acid, the hydrogen is displaced 
from the acid by the zinc, or the zinc is substituted chemically 
for the hydrogen. Similarly, hydrogen is displaced (in 
part) from water by sodium, or sodium is substituted (in 
part) for the hydrogen of water. Look over the equations 
again (§§ 68, 69, 70). 

We define displacement as a chemical change in which 
one element displaces another from a compound. 

Metals differ in their power to displace hydrogen from 
acids or water. Based on their relative displacing power, 
they can be arranged in a hydrogen displacement series : 

Potassium, Sodium, Calcium, Magnesium, Aluminum, Zinc, Iron, 
Lead, Copper, Mercury, Silver, Gold. 

The metals following lead do not displace hydrogen from 
acids or water. That is, we cannot prepare hydrogen from 
acids and the metals copper, mercury, silver, and gold. Nor 
do the metals beyond lead liberate hydrogen from water. 

73. Physical properties of hydrogen. — Hydrogen has no 
taste or color. The pure gas has no odor, though hydrogen 
as ordinarily prepared has a disagreeable odor, due mainly 
to impurities in the metals used. Hydrogen is only very 
slightly soluble in water. 


HYDROGEN 


53 


Hydrogen is the lightest known substance. Volume for 
volume hydrogen is about one fourteenth as heavy as air 
and one sixteenth as heavy as oxygen. One liter at 0° C. 
and 760 mm. weighs only 0.09 gm. (exactly 0.08987 gm.). 
Hydrogen, being so light, diffuses rapidly, ^.e., it quickly 
passes through porous substances {e.g., unglazed porcelain, 
rubber covering of balloons, and thin sheets of metal). If 
a bottle of hydrogen is exposed to the air a minute or two, 
and then a lighted match is dropped in, the match merely 
burns; if hydrogen were present, a loud explosion would 
have occurred. 

74. Chemical properties of hydrogen. — At ordinary 
temperatures and under ordinary conditions, hydrogen is 
not an active element. A 
mixture of hydrogen and 
oxygen can be kept indefi¬ 
nitely at the temperature of 
the laboratory. But if the 
mixture is heated to about 
800° C. or if a fiame is 
brought very near it, the 
gases unite with a violent 
explosion. 

Under special conditions 
hydrogen unites with oxy¬ 
gen quietly. Thus, hydro¬ 
gen burns quietly in oxygen 
or air if the hydrogen is forced through a small opening and 
then cautiously ignited. 

To prove this, hydrogen is generated in the apparatus shown in Fig. 
33; the gas passes through the drying tube and escapes in a fine stream 
through the small opening in the platinum tip. After all the air has 
been driven out of the whole apparatus by the hydrogen, the gas is 
lighted by holding a lighted match at the end of the tip. The flame can 
hardly be seen and is very hot. If a small, dry, cold bottle is held oyer 
the flame, water vapor condenses as a film inside the bottle. That is, 
hydrogen in burning unites with oxygen, just as other burning sub¬ 
stances do. Thus: — 

2 H 2 + O 2 = 2 H 2 O 

Hydrogen Oxygen Water 



Fig. 33. — Apparatus for burning 
hydrogen. 





















54 


A BRIEF COURSE IN CHEMISTRY 


This striking chemical characteristic of hydrogen illus¬ 
trates the chemical change called combination and also 
oxidation (§§ 27, 28). The two elements, hydrogen and 
oxygen, unite to form the compound water; hydrogen is 
oxidized, and the product might be called hydrogen oxide, 
though it is called by its more familiar name water. 

The film of water that may be seen on the bottom 
of a vessel placed over a lighted gas range or a Bunsen 
burner is the condensed vapor formed by the burning 
hydrogen, and hydrogen compounds, in the illuminating 
gas. Organic substances containing hydrogen, such as 
wood and paper, when burned, yield water as one of their 
products. 

Although a small jet of hydrogen when forced out of a 
generator or a tank of the gas burns quietly in air or in 
oxygen, a mixture of hydrogen and air burns so rapidly that 
the combustion is practically an explosion. Therefore the 
air should be fully expelled from the apparatus in which 
hydrogen is being generated and all leaky joints should be 
tightened before the gas is collected; no flames, large or 
small, should be near. Neglect of these precautions has 
caused Serious accidents. 

Certain metals under special conditions hasten the rate 
at which hydrogen combines with elements. Thus, if a 
mixture of hydrogen and oxygen is passed over finely di¬ 
vided platinum, the gases, which ordinarily would scarcely 
combine at all, now begin to react to form water. Similarly, 
hydrogen and nitrogen can be made to unite into the com¬ 
pound ammonia (NHs), if the mixture is passed under proper 
conditions over certain metals. In the World War the Ger¬ 
mans made ammonia this way, and considerable is manu¬ 
factured in the United States by this method. The platinum 
and the other metals, as far as we know now, do not undergo 
a permanent chemical change in these reactions. They 
hasten a very slow chemical reaction. A metal, or any 
other substance, which acts thus is called a catalyst or a 
catalytic agent (§ 31). Its function is sometimes compared 
to that of lubricating oil on a machine. (See also hydro¬ 
genation in § 78.) 


HYDROGEN 


55 


76. Hydrogen burns but it does not support combustion. 

— These facts are illustrated by putting a lighted taper into 
an inverted bottle of hydrogen (Fig. 34). The taper ignites 
the hydrogen, which burns at the mouth of the bottle. The 
taper does not burn inside the bottle, but when it is with¬ 
drawn through the burning hydrogen it is relighted. 

76. Hydrogen is a reducing agent. — Hydrogen not only 
combines energetically with free oxygen, but it also with¬ 
draws oxygen * from compounds. This chemical removal 
of oxygen, as we learned in § 48, is called re¬ 
duction, and the substances that remove the 
oxygen are called reducing agents. Hydrogen 
is an energetic reducing agent. 

When oxides of certain metals are heated 
in a current of hydrogen, the oxygen of the 
oxide is chemically removed and combines 
with the hydrogen to form water; the metal 
is left uncombined. Thus, by heating copper 
oxide in hydrogen, or a gas which contains free 
hydrogen, e.g., illuminating gas, water and 
metallic copper are produced. Chemically 
speaking, the copper oxide is reduced by the 
hydrogen. The chemical change is also an 
example of substitution (the hydrogen being substituted for 
the metal), thus : — 

CuO 4- H 2 = H 2 O + Cu 

Copper Oxide Hydrogen Water Copper 



Fig. 34.— Hy¬ 
drogen burns 
but does not 
support com¬ 
bustion. 


This chemical change can also be interpreted from the 
standpoint of oxidation, because the hydrogen is oxidized 
to water at the same time the copper oxide is reduced. In 
fact, the processes of reduction and oxidation are closely 
related and either one may be emphasized in interpreting 
the chemical change. In its simplest form, reduction is the 
opposite of oxidation. Later we shall see that the terms 
oxidation and reduction are both used in a broader sense 
(§ 317). 

77. Test for hydrogen. — A simple test for hydrogen 
is that it extinguishes a small flame, such as a blazing taper 








56 


A BRIEF COURSE IN CHEMISTRY 


or joss stick, but is lighted at the same time, often with an 
explosion, and continues to burn until the gas is exhausted. 
A conclusive test is that it burns with a colorless, hot flame 
and forms water as the sole product. 

78. Uses of hydrogen. — On account of its extreme light¬ 
ness, hydrogen is used to fill balloons and some dirigible 
airships; small balloons are usually filled with coal gas 
and larger dirigible craft in the United States are filled with 
the non-combustible gas helium (§ 154). 

The intense heat of the hydrogen flame is utilized in the 
oxy-hydrogen blowpipe. The burner consists of two pointed 



Courtesy General Electric Co. 

Fig. 35. — Workman welding a gear casing with a flame of atomic hydrogen. 


metal tubes. The inner and smaller one is for oxygen, and 
the outer and larger one for hydrogen; the gases are obtained 
from tanks and are forced through the small openings of the 
tubes by the pressure maintained in the tanks. The flame 
is used to melt quartz and platinum. The oxy-hydrogen 
flame has been largely replaced by the electric furnace and 
the oxy-acetylene flame (§ 35). (Acetylene is a compound 
of carbon and hydrogen — C 2 H 2 .) 

A special form of hydrogen, called atomic hydrogen, burns 
with such a hot flame it is used in welding (Fig. 35). 

The most extensive use of hydrogen is in transforming 




HYDROGEN 


57 


oils into solid edible fats. When certain oils, e.g., cotton¬ 
seed oil, are mixed with a small quantity of finely divided 
nickel, heated, and treated with hydrogen under pressure, 
the oil and hydrogen unite and form a solid fat. The nickel 
acts as a catalyst, and is ultimately removed by filtering 
(from the melted fat). “ Crisco ” and similar cooking fats 
are made from cotton-seed oil by this process; objectionable 
fish oils are also changed into fats suitable for making soap. 
This chemical change, which consists in the direct chemical 
addition of hydrogen hastened by a catalyst, is called hydro¬ 
genation. 


EXERCISES 

1 . How can hydrogen be distinguished from (a) oxygen, (6) carbon 
monoxide, (c) carbon dioxide, (d) air? 

2. Summarize the main physical properties of hydrogen. 

3. Why is there danger of an explosion in generating hydrogen? 
How can the danger be avoided? 

4. Define and illustrate (a) reduction and (6) reducing agent. Com¬ 
pare with (a) oxidation and (6) oxidizing agent. 

6. State the test for hydrogen. 

6. Name the metals in the order of their displacing power. 


PROBLEMS 

1. Sulfuric acid contains 2.04 per cent of hydrogen. How many 
grams must be decomposed to yield 85 gm. of hydrogen ? 

2. Water contains 11.18 per cent of hydrogen. How many gm. of 
hydrogen can be prepared from 230 gm. of water? How many cc.? 

3. A student prepared enough hydrogen to fill six bottles, each hold¬ 
ing 250 cc. How many grams were prepared? 


SUGGESTIONS FOR LABORATORY WORK 

(References are to NewelFs Laboratory Exercises in Chemistry) 

Exercise *9 — Preparation and Properties of Hydrogen. 
Exercise S8 — Preparation of Hydrogen (Short Method). 
Exercise *10 — Reduction of Copper Oxide by Illuminating Gas. 
Exercise *11 — Reaction between Sodium and Water — T. 
Exercise S9 — Reaction between Zinc and Sulfuric Acid — T. 
Exercise SIO — Reduction of Copper Oxide by Hydrogen — T. 


58 


A BRIEF COURSE IN CHEMISTRY 


SUPPLEMENTARY SECTIONS FROM PART II 

336. Hydrogen and energy. 

336. Hydrogenation of fats. 

The following sections may be selected from Topic VI: — 

337. Gases. 

338. Measuring gases. 

339. Finding the volume of a gas. 

340. Law of Charles. 

341. How we apply the law of Charles. 

342. Law of Boyle. 

343. How we apply the law of Boyle. 

344. How we correct a gas volume for both temperature and pressure. 
346. Finding the weight of a given volume of a gas. 

346. Chemical laws and theories. 

347. The kinetic-molecular theory of the structure of gases. 


CHAPTER VII 


WATER 

79. Distribution of water. — Water is always present 
in the air as a vapor, which is continually condensing into 
rain, clouds, mist, fog, dew, hail, frost, or snow. Water 
occurs as a liquid in vast quantities on and beneath the 
surface of the earth. As a solid (snow and ice) it covers the 
tops of high mountains and the polar regions of the globe. 
Many common foods contain water, often a large per cent, 
e.g., milk 87, potatoes 78, eggs 73, beefsteak 62. All living 
things contain water, often to an unusual extent. Thus, 
the human body is nearly 70 per cent water. 

80. Impurities in water. — Water is never found pure in 
nature. Even rain water contains gases and dust washed 
from the air. When rain strikes the ground, it begins at 
once to take up impurities from the soil and rocks. Some of 
the water flows along the surface, becoming more and more 
impure, and Anally reaches the ocean. Some of it soaks 
into the ground and percolates through the soil. On its 
journey underground the water dissolves mineral matter 
and gases. The mineral matter is usually common salt and 
compounds of calcium, magnesium, and iron; the most 
common gas is carbon dioxide. If the amount of dissolved 
matter is so unusual as to give the water a marked taste or 
medicinal properties, the water is called mineral water. 
Many common mineral waters have limited medicinal value. 

Water containing calcium and magnesium compounds 
is called hard water, but in soft water, such as rain water, 
these compounds are absent. River water contains earthy 
impurities brought by the underground and surface water; 

59 


60 


A BRIEF COURSE IN CHEMISTRY 


it is also often contaminated with compounds formed by the 
decomposition of animal and vegetable matter (called organic 
matter), refuse from manufactories, sewage, and harmful 
bacteria. Ocean water contains about 3.5 per cent of dis¬ 
solved mineral substances, largely common salt and com¬ 
pounds of calcium and magnesium. The peculiar taste of 
ocean water is due chiefly to these substances. 

Besides dissolved mineral matter, water may contain 
suspended matter, such as flne particles of clay or sand and 
partly disintegrated organic substances. 



Fig. 36. — Purification of water by spraying it into the air. An aerator 
of the New York City water system by which 376 million gallons are 
purified daily. 


81. Purification of water. — It is often necessary to remove 
certain impurities from water used for drinking and for in¬ 
dustries. Thus, water for domestic use should be free from 
poisons, iron compounds, and as much as possible, magnesium 
and calcium compounds. And water for steam boilers 
should be free from calcium and magnesium compounds to 
prevent the formation of boiler scale (§ 298). 

The water used in some towns and cities is purified by 
filtering it slowly on a large scale through layers of sand and 
gravel. By such a filter, suspended matter is almost com¬ 
pletely removed, and certain impurities are decomposed 
by microorganisms in the upper layer of the filter. A sand 




WATER 


61 


filter must be frequently cleaned. In certain localities the 
water is stored in a large settling basin or reservoir; here the 
suspended solid matter slowly settles, the process sometimes 
being hastened or aided by adding aluminum sulfate, which 
forms a sticky substance to which fine particles cling. In 
other localities the water is freed from organic matter by 
spraying it into the air (Fig. 36). 

The most effective method of destroying organic matter 
(especially bacteria) in water is by chemical treatment. 
For this purpose chlorine is used. A small amount of liquid 



Fig. 37. — Purification of water by chlorine. A chlorinator of the New 
York City water system by which 400 million gallons are purified daily. 


chlorine (about 20 pounds to 1,000,000 gallons of water) is 
allowed to flow into the entering supply at the pumping 
station (Fig. 37). The water in a swimming pool can be 
purified by chlorine: the proportions are about 0.5 pound 
of chlorine (or 1.5 pounds of bleaching powder) to 100,000 
gallons of water. 

Water containing bacteria can be purified on a small 
scale or in an emergency, e.g.,, for household use during an 
epidemic, by boiling the water ten or fifteen minutes, and 
then putting it in a bottle or a jar stoppered with cotton. 
The bottle or jar for keeping the water should be boiled in 












62 


A BRIEF COURSE IN CHEMISTRY 


water before use. The small filters sometimes used on faucets 
in the house are not reliable purifiers because they accumulate 
dirt through which the water must pass. 

If the purity of a drinking water is doubtful, a sample 
should be subjected to a chemical and microscopic examina¬ 
tion, supplemented by a rigid sanitary inspection of the sur¬ 
roundings and the source of supply. 

82. Distillation. — Water can be purified by distillation, 
^.e., by boiling the water and condensing the steam. By 
this process the water is separated from suspended and non- 



Fig. 38. — Liebig condenser arranged for distillation of water. The vapor 
from the water boiling in the flask A condenses in the inner tube B, 
owing to the decrease in temperature caused by the current of cold water 
in the outer tube CD. The condensed water, called the distillate, drops 
off the lower end into the receiver E, while the non-volatile impurities 
remain behind in the flask. 


volatile impurities, e.g., mineral matter and most organic 
matter. Water thus purified is called distilled water. 

Distillation is performed in a condenser*. One form, 
known as the Liebig condenser, is often used in the laboratory 
(Fig. 38). A simplified distilling apparatus for laboratory 
use is shown in Fig. 39. 

Other forms of apparatus are used, especially for continuous 
work. One is shown in Figure 40. Cold water enters the con¬ 
denser C at the lower inlet F, and is kept level in the chamber 
A by the upper outlet G. The chamber A is heated by the 
burner E, and steam in passing through B and down through 
the inner tube of the condenser C drops off at D as dis- 

















WATER 


63 


Fig. 39. — Simplified distilling appa¬ 
ratus for use in the laboratory. 
The water vapor from A condenses 
in B, which is cooled by the water 
in C. 


tilled water. Distilled water 
is prepared on a large scale by 
boiling water in a metal ves¬ 
sel and condensing the vapor 
in a block tin or quartz pipe 
coiled around the inside of a 
vessel through which a current 
of cold water is flowing. Glass 
is not used, because water 
dissolves traces of substances 
from the glass and these 
would contaminate the dis¬ 
tilled water. 

Distilled water has a flat 
taste. If it is to be used in 
making beverages, it is aer¬ 
ated by bubbling air through 
it; the water soon acquires 
the accustomed flavor. Much 

distilled water is used in the chemical laboratory to prepare 

solutions and in experiments 
requiring water free from or¬ 
ganic and mineral matter. 
Distilled water is used to pre¬ 
pare medicines which require 
water free from mineral and 
organic matter; it is also used 
in the laboratory to prepare 
certain solutions, e.g., silver or 
lead compounds. 

83. Some physical proper¬ 
ties of water. — At ordinary 
temperatures pure water is a 
tasteless and odorless liquid. 
It is usually colorless, but 
thick layers are blue. It is a 
liquid between 0° and 100° C. 
„ ^ , j. ..M,- This means that it freezes 

water. at 0 C. and boils at 100° C. 










































64 


A BRIEF COURSE IN CHEMISTRY 


Below the freezing point it is a solid; above the boiling 
point it is a gas. 

Most liquids expand when heated and contract when 
cooled. But water behaves exceptionally with changes in 
temperature. If water at 100° C. is gradually cooled, it 
contracts until 4° C. is reached; if the cooling continues, it 
expands until 0° C. is reached. Hence at 4° C. a given 
volume contains the greatest weight of water, i.e., at 4° C. 
water has its maximum density (Fig. 41). 

The density of a liquid is the weight of 1 cc. of the liquid. 
In the case of water, 1 cc. weighs 0.95838 gm. at 100° C., 
1 gm. at 4° C., and 0.99987 gm. at 0° C. Since water has 
its maximum density at 4° C., the density of water at 4° C. 

100 ° _^_ 4 ° ^ 0 ° 

Contracts Expands 


Expands Contracts 

loo^ ^ 4^ ^ 0^ 

Fig. 41.—Water has the maximum density at 4° C. In cooling, it con¬ 
tracts between 100° and 4° C., but expands between 4° and 0° C. Hence 
a given volume weighs the most at 4° C. 


is taken as 1. Water at this temperature is the standard 
for defining the specific gravity of solids and liquids. By 
specific gravity we mean the weight of the solid or liquid 
compared with the weight of an equal volume of water. 
Thus, the specific gravity of gold is 19, i.e., a piece of gold is 
19 times heavier than an equal volume of water at 4° C. 

The expansion of water when cooled from 4° to 0° C. is 
very slight, but the change is exceedingly important in nature. 
When the water on the surface of a lake or river cools, it 
contracts, and since it is heavier (volume for volume) than 
the warmer water beneath, it sinks. The warmer water 
rises, becomes cool, and likewise sinks, thus causing a cir¬ 
culation which continues until all the water from the surface 
to the bottom has a temperature of 4° C. If the cooling 
continues, the surface water expands and remains on the 
top, because it is lighter than the water beneath. Hence 







WATER 


65 


when the temperature of the air falls to 0° C., this upper 
layer of water freezes and protects the remaining water from 
the cold air, thus stopping the circulation. Should the 
circulation continue, as the temperature fell from 4° to 0° C., 
the whole body of water would finally freeze from top to 
bottom. This condition would not only destroy plants and 
animals in the water, but would also profoundly affect our 
climate by the accumulation and slow melting of the ice. 

The temperature at which water solidifies or freezes is 
0° C. (or 32° F. — Fahrenheit). This temperature (0° C.) 
is popularly called the freezing point of water. Strictly 
speaking, the freezing point of water is not the temperature 
at which ice forms, but the temperature of water in which 
ice is floating; this temperature is 0° C. and stays so until 
the ice is melted. When water freezes, it expands nearly 
one-tenth of its volume. More exactly, 100 cc. of water 
produce 109 cc. of ice. The specific gravity of ice is very 
nearly 0.92. Hence ice floats in water. 

The pressure exerted by water when it expands in freezing 
is powerful. Vessels or pipes completely filled with water 
often burst when the water freezes. It is a popular idea 
that “ thawing out ” a pipe bursts it. This is not true, 
because ice contracts when it melts. As a matter of fact, 
pipes crack as soon as the water freezes, and later when the 
ice melts, the water flows out of the crack. 

When water is heated sufficiently in an open vessel, the 
liquid becomes a vapor and escapes rapidly. When the 
temperature of the water reaches 100° C. (or 212° F.), the 
water boils, ^'.e., it changes rapidly into vapor and the tem¬ 
perature stays at 100° C. This vapor is invisible, but as it 
leaves the vessel it cools and condenses quickly into a cloud 
of minute drops of liquid water. This cloud is called 
steam. 

84. Evaporation. — Liquid water is always changing into 
water vapor. This process is called evaporation. Water 
vapor is escaping constantly on an enormous scale from the 
surface of the ocean, lakes, and rivers. On a small scale 
water vapor is escaping from puddles of water, moist soil, 
and wet objects. We speak of this process as drying. 


66 


A BRIEF COURSE IN CHEMISTRY 


85. Chemical properties of water. — We have already 
seen that water at ordinary temperatures interacts with 
sodium and potassium, and at higher temperatures with 
iron (§ 69). 

Water combines directly with many oxides. Thus, 
calcium oxide (CaO) combines with water and forms calcium 
hydroxide (Ca(OH) 2 ). Similarly, sulfur dioxide (SO 2 ) forms 
the compound sulfurous acid (H 2 SO 3 ). Such oxides are 
often called anhydrides (see § 219). 

Water combines with certain substances and thereby forms 
compounds called hydrates, especially when such substances 
separate from a solution by crystallization. Thus, from a 
solution of copper sulfate the blue crystals obtained are a 
hydrate having the formula CUSO 4 .5 H 2 O. The 5 H 2 O in 
this formula is separated by a period from the CUSO 4 to 
indicate that the crystallized solid is a hydrate. 

86. Hydrates and water of hydration. — Many familiar 
substances are hydrates, e.g., washing soda is a hydrate of 
sodium carbonate and alum of aluminum potassium sul¬ 
fate. All hydrates, when heated, give up water and form the 
corresponding anhydrous compound. Thus, the familiar 
blue solid called copper sulfate is a hydrate (CUSO 4 .5 H 2 O) 
and the gray-white powder formed by heating it is the corre¬ 
sponding anhydrous compound (CUSO4). Hydrates are 
made up of two parts, water and the anhydrous substance. 
The proportion of water of hydration differs in different 
hydrates. But each hydrate contains an unvar 3 dng per 
cent of water, which is given off on sufficient heating, e.g.y 
hydrate of copper sulfate loses 36.36 per cent. Most hy¬ 
drates are crystals, often colored, whereas anhydrous sub¬ 
stances are usually white, or slightly colored, powders. 
Anhydrous substances often become readily hydrated again. 
Thus, when water is poured upon the gray anhydrous copper 
sulfate, the solid turns blue, and with sufficient water a blue 
solution is obtained, from which blue crystals of hydrated 
copper sulfate settle out again. 

Some hydrates are so unstable they lose water merely 
by exposure to air, crumble to a lusterless powder, and, if 
colored, lose their color. This property is called efflores- 


WATER 


67 


cence, and such substances are said to be efflorescent or to 
effloresce. Washing soda, alum, sodium sulfate, and borax 
are examples of hydrates which effloresce readily. 

87. Water as a solvent. — Water dissolves a great many 
substances, ^.e., they disappear when put into water. The 
liquid which results from this process of dissolving is called 
a solution. The dissolved substance is called the solute, 
and the water the solvent. A water solution of a substance 
is sometimes called an aqueous solution — the kind we are 
considering in this chapter. 

The degree to which substances dissolve is sometimes 
called their solubility (compare § 90). A solution which 
contains a small proportion of solute is called dilute ; one 
containing a large proportion is called concentrated. Addi¬ 
tion of water to a solution is called diluting. The partial 
removal of water by evaporating is called concentrating or 
evaporating to crystallization; the total removal of water is 
called evaporating to dryness. 

Sometimes the word concentration is used to describe the 
degree of solubility. If the concentration reaches a max¬ 
imum at a specified temperature the solution is called 
saturated (compare § 92). 

Substances differ widely in their solubility. Some are 
very soluble, e.g., ammonia (gas), alcohol (liquid), sugar 
(solid). Others are only very slightly soluble, e.g.^ hydrogen 
and oxygen (gases), gasolene and oils (liquids), and sand and 
limestone (solids). The substances mentioned in the last 
sentence, and others like them, are often called insoluble 
because the amount dissolved is insignificant. A large class 
of substances may be described as moderately soluble. 

88. Solutions of gases. — Some gases, like ammonia, 
are very soluble, while others, such as oxygen and nitrogen, 
are only slightly soluble. As a rule, the solubility of a gas 
decreases with rise of temperature. Pressure, too, infiuences 
the solubility of gases. Thus, as we have already seen (§ 56), 
carbon dioxide is forced into cylinders of water in preparing 
soda water. When the pressure is decreased by opening 
the valve, the gas escapes rapidly and causes the soda water 
to froth or foam. Bubbles caused by escaping carbon di- 


68 


A BRIEF COURSE IN CHEMISTRY 


oxide may also be seen when the cap is removed from a 
bottle containing a charged beverage (Figs. 21, 22). 

This rapid escape of a gas is called effervescence. Under¬ 
ground waters often contain considerable carbon dioxide, 
owing to the great pressure to which subterranean gases are 
subjected. Hence, many natural mineral waters effervesce 
when they come to the surface. 

89. Solutions of liquids. — Some liquids, such as sulfuric 
acid, alcohol, and glycerin, dissolve in water in all propor¬ 
tions; others, such as gasolene and kerosene, are very 
slightly soluble. 

Certain liquids dissolve to a limited extent in water. Ether 
is an example. 

90. Solutions of solids. — The solubility of solids in water 
depends on the substance. Some, like sand, dissolve scarcely 
at all and are often described as insoluble. Others, like salt, 
are moderately soluble, while still others, like magnesium 
chloride or sodium hydroxide, are highly soluble. 

The degree of solubility also depends on temperature. 
Solubility in most cases increases with rise of temperature. 
A few solids {e.g.^ calcium hydroxide) are less soluble in hot 
water than in cold, and a few others {e.g., sodium chloride) 
dissolve to about the same degree in hot and cold water. 

There is a limit to the solubility of most substances. As a 
rule, a given weight of water at a fixed temperature will 
dissolve only a definite weight of solid; and this is the case, 
even though more undissolved solid is available for solution. 
A solution conforming to the conditions just stated is said 
to be saturated. For example, at 20° C. 100 gm. of water 
dissolve 7.2 gm. of potassium chlorate. This solution is 
saturated, for if more potassium chlorate is added, it remains 
undissolved. A saturated solution may also be defined as 
a solution in which maximum concentration is maintained in 
contact with some of the undissolved solid. 

91. How solubility is expressed. — For ordinary pur¬ 
poses, solubility may be expressed by such general terms as 
insoluble, slightly solublej or very soluble. A more accurate 
way is to state the number of grams of solid dissolved by 
100 gm. of water at a fixed temperature, as shown in Table II. 


WATER 


69 


TABLE II. — Solubility of Solids in Water 


Solids 


Number of Grams in Solution in 
100 Grams of Water 



10° C. 

20° C. 

100° C. 

Ammonium chloride. 

33.3 

37.2 

77.3 

Calcium hydroxide. 

0.176 

0.165 

0.077 

Potassium chloride. 

31.0 

34.0 

56.7 

Potassium nitrate. 

20.9 

32.0 

246.0 

Sodium chloride. 

35.8 

36.0 

39.0 


A useful way of representing solubility is by a solubility 
curve. That is, by a curve drawn through a series of points 
on cross-section paper, located by finding experimentally 
the number of grams of the substance soluble in 100 gm. of 
water at many temperatures. The curves of several sub¬ 
stances are shown in Fig. 42. The temperature is read along 
the vertical lines and the number of grams of solute in 
100 gm. of water along the- horizontal lines. 

Many facts can be read from solubility curves. For 
example, (1) some substances are very soluble and others 
only moderately so, (2) the solubility of some substances 
increases rapidly with rise of temperature, and (3) the 
increase of solubility of certain substances is slow or slight. 
So also, if we wish to know the temperature at which a 
certain number of grams, say 40 gm. of potassium chlorate, 
are held in solution by 100 gm. of water, we find where the 
horizontal line numbered 40 cuts the potassium chlorate 
curve, and then follow the vertical line down to the tempera¬ 
ture number, where 80° C. is found. 

92. Solution, crystals, and crystallization. — If hot con¬ 
centrated solutions are cooled, or if enough of the solvent is 
removed from a cold solution by evaporation, part of the 
solute.separates from the solvent in the solid form, often in 
beautiful shapes called crystals. The process of separating 
a dissolved solid from its solution by either of these ways is 
called crystallization. The shape and color of the crystals 
















70 


A BRIEF COURSE IN CHEMISTRY 


are characteristic of the particular substance and serve to 
identify it. Thus, common salt crystallizes in white cubes. 

93. Supersaturated solution. — Crystals are not always 
deposited from a cooled or a concentrated solution, as just 



stated. Thus, a hot, very concentrated solution of some 
solids, such as sodium sulfate or sodium thiosulfate, deposits 
no crystals when the clear solution cools, although the solu¬ 
tion actually contains more solute than the solvent could 
ordinarily dissolve at the lower temperature! Solutions 
which contain more solute than is needed for normal satura- 




















































WATER 


71 


tion are called supersaturated. Supersaturation occurs 
only under special conditions. For example, if a super¬ 
saturated solution is stirred or violently shaken, crystals 
begin to form. Moreover, if a fragment of the solid is 
dropped into the supersaturated solution, crystals very 
quickly form upon the fragment and soon accumulate in a 
conspicuously large quantity. 

94. Deliquescence. — Many substances when exposed to 
air become moist, and sometimes even dissolve in the ab¬ 
sorbed water. This property is called deliquescence and 
the substances are said to deliquesce, or to be deliquescent. 



Fig. 43. Apparatus used to dry substances. On the right is a desiccator 
containing calcium chloride (in the bottom) which dries the substance in 
the crucible and the air in the desiccator. The other two are used to dry 
gases. 

Deliquescence is a property of very soluble substances. 
Calcium chloride, potassium carbonate, zinc chloride, sodium 
hydroxide, and potassium hydroxide belong to this class. 
Common salt often deliquesces, especially in damp weather, 
owing to small quantities of magnesium and calcium chlo¬ 
rides which are present as impurities. The property of 
deliquescence is utilized in the laboratory to dry substances, 
calcium chloride often being employed for this purpose. 
Different forms of apparatus are used (Fig. 43). 

95. What is a solution ? — Let us try three experiments. 
(1) If we shake a little salt with water, the salt dissolves. 
We have a solution, Le., a mixture in which the salt is uni¬ 
formly distributed and from which the salt will not settle. 


















72 


A BRIEF COURSE IN CHEMISTRY 


(2) If we shake a little powdered starch with cold water, 
the starch is distributed more or less uniformly throughout 
the liquid. But on standing, the starch begins to settle, 
until finally the mixture separates into water and starch. 
The starch was not dissolved, but merely suspended in the 
water. Such a mixture of a solid and water is called a 
suspension. 

(3) If we shake kerosene vigorously with water, the oil 
breaks up into fine drops which are distributed throughout 
the liquid. But after a time the oil separates from the 
water. Such a mixture of a liquid and water is called an 

emulsion. 

Suspensions and emulsions have a common property, 
viz., the suspended substance in time separates from the 
water. On the other hand, the dissolved substance does not 
separate from a solution. 

This distinctive property of a solution, viz., non-settling 
of the solute, is doubtless due to the fact that in a solution 
the dissolved particles are exceedingly minute — too minute 
to be seen through a microscope or detected by a beam of 
light. If a beam of light is passed into a dark room, the 
path is revealed by the suspended dust particles that reflect 
the light; this effect may be seen when a sunbeam comes 
through an opening in a blind or a hole in a curtain. But 
when a strong beam of light is passed through a solution, no 
bright path is revealed, because the dissolved particles are 
very much smaller than dust particles, indeed too small to 
reflect light. 

Certain substances, however, form clear mixtures with 
water from which the substance does not settle readily, if at 
all, nor can it be removed by ordinary filtering. Thesn 
liquids look homogeneous, i.e., they look just alike through¬ 
out. And yet, if we pass a beam of light through them, the 
path is bright, thereby proving that these mixtures contain 
particles in suspension. Such mixtures are not true solutions, 
but colloidal solutions, i.e., mixtures in which the particles 
in suspension are very fine — almost as fine as the particles 
in a true solution. Colloidal solutions should really be 
called colloidal suspensions, but the term colloidal solution 


WATER 


73 


is convenient. Many substances can be reduced to the 
colloidal state. Such substances are called colloids. Typical 
examples are starch, clay, silicic acid, and metals like gold. 
(See §§ 261, 446, 498, 511.) 

In true solutions, then, the particles of the solute are in an 
exceedingly fine state of division. Whereas in suspensions 
and emulsions the particles are much larger — large enough 
to be seen through a microscope and often with the eye. 
Between these two classes come colloidal solutions; in them 
the particles are too fine to settle out, though large enough 
to reflect light, ranging in size from those 
in true solutions to those in typical sus¬ 
pensions. 

In passing from true solutions through 
colloidal solutions to ordinary suspensions 
(and emulsions), the change is gradual, 
not abrupt, since the distinction is based 
fundamentally on the size of the particles. 

When a colloidal solution is examined 
through an ultramicroscope (a very pow¬ 
erful microscope arranged for sending 
a strong light horizontally through the 
liquid), tiny bright specks are seen dart¬ 
ing about in a zigzag motion. These 
specks are colloidal particles. 

The (enlarged) path of one of these 
points seen m a colloidal solution is com- particle, 
plicated (Fig. 44). This motion is called 
the Brownian movement from Robert Brown, the Scotch 
botanist who first studied it. It is seen in all colloidal 
solutions regardless of the nature of the particles or the 
age of the solution. The erratic movement of the colloidal 
particles is due to the ceaseless, irregular bombardment of 
the particles by the molecules of water. The movements 
and the size of the colloidal particles have been carefully 
investigated, and the results furnish conclusive evidence 
that the characteristic particles in a true solution are mole¬ 
cules and in a colloidal solution are groups, or aggregates, 
much larger than molecules. 





74 


A BRIEF COURSE IN CHEMISTRY 


96. Composition of water. — Water is a compound of 
hydrogen and oxygen. That is, its constituents are the 
elements hydrogen and oxygen, and they are chemically 
combined in a constant ratio. 


We have already learned that hydrogen is one constituent 
of water. (1) Hydrogen is liberated when certain metals, 
e.g., sodium, potassium, and iron, interact with water (§69). 
(2) Hydrogen is one of the gases liberated when an electric 
current is passed through an acid (or an 
alkaline) solution of water (§§ 22, 69). 
(3) Hydrogen if burned forms water (H 2 O) 
(§ 76). 

Similar facts have shown us that oxygen 
is a constituent of water. (1) Oxygen com¬ 
pounds are formed when metals react with 
water, e.g., iron and steam form iron oxide 
and sodium forms sodium hydroxide (§ 69). 
(2) Oxygen is also liberated when an electric 
current is passed through water containing 
an acid (or an alkali) (§ 22). (3) When 

hydrogen is burned in air or oxygen, water 
(H 2 O) is formed (§ 75). (4) Oxygen is 

formed when water and chlorine interact. 








Fig. 45. — Experi¬ 
ment to show 
that oxygen is 
one constituent 
of water. 


We can show (4) by a simple experiment. A tube 
(Fig. 45) about a meter long and closed at one end 
is completely filled with chlorine water (prepared 
by saturating water with chlorine) (§ 124 .) The 
open end is immersed in a vessel containing some 
of the same solution, and the whole apparatus is 
placed in the direct sunlight. Tjie chlorine and 
water interact. Bubbles of gas soon appear in the 
liquid, and after a few hours a small volume of gas collects at the top 
of the tube (Fig. 45). The gas can be shown to be oxygen by the usual 
test, viz., rehghting a glowing joss stick or a spHnt of wood. 


97. Electrolysis of water. — Pure water does not conduct 
electricity, but a solution containing an acid, a base {e.g., 
sodium hydroxide), or a salt {e.g., sodium chloride) does 
conduct electricity. The decomposition of such a water 
solution by electricity, called traditionally the electrolysis 


















WATER 


75 


of water, shows by a single experiment that water consists 
of hydrogen and oxygen. 

It can be done in a special form of apparatus called the 
Hofmann apparatus (Fig. 46). 


A mixture of water (10 vols.) and concentrated sulphuric acid (1 vol.) 
is pomed into the apparatus until the reservoir is half full (after the 
stopcocks have been closed). 

The electric current may be obtained from a 6-volt storage battery, 
or four dry cells connected in series, or a direct street current reduced by 
two 50-watt lamps. The wires are connected with 
the piece of platinum near the bottom of each tube, 
and as soon as the current passes, bubbles of gas ap¬ 
pear on the platinum, rise, collect in the upper part 
of the tubes, and slowly force liquid from each tqbe 
up into the reservoir. 

The current should be shut off when the tube con¬ 
taining the larger volume of gas is about three-fourths 
full. Assuming that the tubes have the same diam¬ 
eter, the gas volumes are in the same ratio as their 
heights, which will be found by measurement to be 
approximately two to one. Tests applied to each 
gas (by letting a little out through the stopcock) 
show that the gas having the larger volume is hy¬ 
drogen and that the other gas is oxygen. 


98. How the exact composition of water 
is found. — The experiments just cited and 
described show the qualitative composition of 
water. That is, they show that water is a 
compound of the two elements hydrogen and 
oxygen. But they give us no information 
about the proportion of the elements in the 
compound. To find the quantitative com¬ 
position of water, we must study the results 
of experiments performed for the purpose of 
determining the exact proportions — “ the quantity ” — 
which the two elements combine to form the compound. 


Fig. 46. — Elec¬ 
trolysis of water 
(containing sul¬ 
furic acid). 


in 


Exceptionally accurate determinations of the composition 
of water were completed by the American chemist Morley 
in 1895 (Fig. 47) after twelve years of labor. In his experi¬ 
ments he not only measured the hydrogen and oxygen that 
combined but he also weighed them and the water formed by 
















76 


A BRIEF COURSE IN CHEMISTRY 



combining them. Morley found (1) 1 part by weight of 
hydrogen combines with 7.9395 parts by weight of oxygen 

and (2) hydrogen and 
oxygen unite in the ratio 
of 2.00268 to 1 by vol¬ 
ume. 

99. Dumas* determination 
of the gravimetric composition 
of water. — Another method 
was used by the French 
chemist Dumas in 1843. It 
consisted in passing dry hy¬ 
drogen over heated copper 
oxide. We have seen (§76) 
that in this chemical change 
hydrogen reduces copper oxide 
and thereby forms copper and 
water. If the copper oxide 
and copper are weighed, the 
loss is the weight of the oxy¬ 
gen used. If the water is 
collected and weighed, the 
difference between the weights 
of the water formed and the 
Fig. 47. —The American chemist. Morley oxygen used is the weight of 
(1838-1923), who made a very accurate the hydrogen. The result 
study of the composition of water. obtained by Dumas while ac¬ 

curate for those days, now 
has only historical interest. However, the method can be readily 
duplicated in our laboratories. 


100. Summary of the composition of water. — Experi¬ 
ments show that water consists of the two elements hydrogen 
and oxygen combined in a fixed ratio by weight, viz., 1 to 
7.9395; they are also combined in the ratio of 2.00268 to 
1 by volume. Usually these ratios are stated approximately 
as 2 to 16 by weight and 2 to 1 by volume. Often the 
gravimetric composition of water is stated in per cent, 
viz., 11.18 per cent of hydrogen and 88.82 per cent of 
oxygen. Sometimes, we say, briefly, water is ^ hydrogen 
and f oxygen. 




WATER 


77 


EXERCISES 

1. State three ways by which drinking water can be purified. 

2 . Describe distillation. Sketch a Liebig condenser. 

3. State four important physical properties of water. 

4. As in Exercise 3, chemical properties. 

6. Define and illustrate (a) water of hydration, (6) eflBorescence, 
(c) anhydrous, (d) hydrate. 

6 . Define and illustrate (a) solution, (6) solvent, (c) solute, (d) sat¬ 
urated solution. 

7. Define and illustrate deliquescence. How does it differ from 
eflSorescence and effervescence ? What common substances (a) eflloresce 
and (6) dehquesce? 

8 . At what temperature does water (a) boil and (6) freeze ? 

9. Describe experiments which show that hydrogen and oxygen 
are components of water. 

10. What is the composition of water by weight and by volume? 

PROBLEMS 

1. If 5 gm. of crystallized aluminum sulfate lose 2.43 gm. on heating, 
what per cent of water of hydration does this crystallized substance 
contain ? 

2. By the use of the solubility curves in Fig. 42 answer the follow¬ 

ing : (a) How many gm. of sodium chloride are in solution at 20°, 30°, 
55°, 65°, 0°, 100°? (b) At what temperatures are 60 gm. and 95 gm. 

of potassium bromide in solution ? (c) How much sodium nitrate is in 

solution at 20°, 25°, 30°? (d) As in (c), how much potassium bromide? 

3. How much (a) oxygen and (b) hydrogen could be obtained from 
25 gm. of water? 

SUGGESTIONS FOR LABORATORY WORK 

(References are to Newell’s Laboratory Exercises in Chemistry) 

Exercise 12 — Purifying Water by Distillation. 

Exercise S 11 — Preparation and Properties of Distilled Water — T. 

Exercise *13 — Suspension and Solution of Solids. 

Exercise *14 — Effect of Heat on the Solubility of Solids. 

Exercise *15 — Formation of Crystals. 

Exercise *16 — Effect of Shape on the Solubility of a Solid — T. 

Exercise *17 — Water of Hydration. 

Exercise S 13 — Anhydrous Compounds. 

Exercise *18 — Per Cent of Water of Hydration. 

Exercise 19 — Electrolysis of Water (Acid Solution) — T. 

Note: Exercises S 12 — T, S 14 — T to S 17 — T may be done at 
this time or postponed till the corresponding subjects are studied in 
Topic VIII, Part II. 


78 


A BRIEF COURSE IN CHEMISTRY 


SUPPLEMENTARY SECTIONS FROM PART II 

348. Vapor pressure. 

349. Efflorescence and vapor pressure. 

360. Deliquescence and vapor pressure. 

361. Volumetric composition of water. 

362. Gay-Lussac’s law of gas volumes. 


CHAPTER VIII 


SYMBOLS AND FORMULAS 



101. What S5nnbols mean. — Each chemical element is 
designated by a symbol (§ 14), which is an abbreviation of 
its name. Thus, O is the 
symbol of oxygen, C of 
carbon, H of hydrogen. 

The symbols of the ele¬ 
ments are given in the 
Table on the back inside 
cover. These symbols 
were introduced and first 
used extensively by the 
Swedish chemist Berze¬ 
lius (Fig. 48) about 1811. 

He analyzed many sub¬ 
stances and used the 
symbols to show the 
composition of com¬ 
pounds and to express 
the relations of elements 
and compounds in chem¬ 
ical changes. 

Each symbol also 
stands for one atom of 
an element. Thus, H means one atom of hydrogen. More 
than one free atom is designated by writing the proper 
numeral before the symbol. Thus, 2 H means 2 free atoms 
of hydrogen. To represent atoms in chemical combination, 
a subscript is used. Thus, H 2 means 2 combined atoms of 


Fig. 48. — The Swedish chemist Berzelius 
(1779-1848), who introduced and first 
used the symbols of the elements. 


79 







80 


A BRIEF COURSE IN CHEMISTRY 


hydrogen, as in H2O, H2SO4. Similarly, NH3 means 3 atoms 
of hydrogen combined with 1 atom of nitrogen. 

Each symbol also expresses the atomic weight of the 
element. Thus, 0 represents one atom of oxygen which has 
the atomic weight 16; similarly, H = 1 , C = 12, N = 14 
(see Table on inside of back cover). 

To sum up, a symbol has these meanings : (1) an element, 
(2) one atom of an element, (3) the atomic weight of an 
element. 

102. What are chemical formulas ? — A formula is a group 
of symbols which expresses the composition of a compound 
(§§11 (3), 15). Thus, H 2 O is the formula of water, CuO 
of copper oxide, NH3 of ammonia. In writing a formula, the 
correct number of symbols of the atoms making up a molecule 
of the compound is placed side by side. Thus, CO is the 
formula of carbon monoxide, because 1 molecule of this 
compound consists of 1 atom each of carbon and oxygen. 
Whereas, CO 2 is the formula of carbon dioxide, because 1 
molecule is composed of 1 atom of carbon and 2 atoms of 
oxygen. The S 3 rmbols making up a formula might be 
written in different orders, but usage has determined the 
order in most cases. 

A formula represents one molecule. To designate several 
molecules, we place the proper numeral before the formula. 
Thus, KCIO3 means 1 molecule and 2 KCIO3 means 2 mole¬ 
cules of potassium chlorate. 

In some compounds certain atoms in a molecule act as a 
group in chemical changes. This fact is expressed by in¬ 
closing the group in a parenthesis, e.g., calcium hydroxide 
has the formula Ca(OH )2 because the group OH usually 
acts as a unit. Such groups are often called radicals. Thus, 
OH is the hydroxyl radical, SO 4 the sulfate radical, and NH 4 
the ammonium radical. Sometimes the parenthesis is re¬ 
placed by a period, e.g., C2H5.OH (ethyl hydroxide) and 
CUSO 4 .5 H 2 O (copper sulfate pentahydrate (§ 86)). The 
period and parenthesis are usually omitted from the 
formulas of familiar compounds, e.g., ammonium hydroxide, 
NH 4 OH. 

In some formulas a group is placed inside a parenthesis. 


SYMBOLS AND FORMULAS 


81 


e.g.j Pb(N03)2. This means that the group NO3, in this 
case, is to be multiplied by 2. The expression 2 Pb(N 03)2 
means that the whole formula must be multiplied by 2. 
That is, in 2 molecules of lead nitrate there are 2 atoms of 
lead, 4 of nitrogen, and 12 of oxygen. 

103. Relation between formulas and molecular weights. 

— Since a symbol stands for the atomic weight of an element, 
a formula stands for the sum of the atomic weights repre¬ 
sented by the group of symbols. This sum is called the 
molecular weight. In a few words, a symbol stands for an 
atomic weight and a formula stands for a molecular weight. 
Thus, the symbols H and Cl stand for the atomic weights 
1 and 35.5 respectively, and the formula HCl stands for 
their sum 1 + 35.5, or 36.5. If we know the formula of a 
compound, a simple way of finding the molecular weight is 
to add the atomic weights corresponding to the atoms in the 
formula. Using approximate values, the molecular weight 
of water (H 2 O) is 2 -|- 16 = 18. Similarly, the molecular 
weight of lead nitrate (Pb(N 03 ) 2 ) is 207 -j- 2(14 48) = 

331. Atomic weights are given in the Table on the inside 
back cover of this book. 

104. Relation between the formula and the percentage 
composition of a compound. — The composition of a com¬ 
pound can be expressed by a formula and also in per cent 
(§ 11). The “ composition in per cent ” is often called 

percentage composition.'^ Thus, we can express the 
composition of water by the formula H 2 O, and also by hy¬ 
drogen = 11.18 per cent and oxygen = 88.82 per cent. 

How are formulas and per cents related? The answer is 
simple. The symbols in the formula stand for numbers, 
and the composition in per cent is the mathematical equiva¬ 
lent of the chemical formula. 

The formula of potassium chlorate is KCIO3. Suppose 
we wish to express the composition in per cent. The process 
consists in transposing the chemical formula KCIO3 into the 
equivalent mathematical expression. There are three 
steps: — 

(1) Find the molecular weight by adding the atomic 
weights. 


82 


A BRIEF COURSE IN CHEMISTRY 


(2) Divide the weight of each element by the total molec¬ 
ular weight, 

(3) Multiply each quotient by 100. 

Proceeding as above, we have (1) K = 39, Cl = 35.5, 
and 3 O = 48 (^.e., 3 X 16), and the molecular weight is 
their sum, 122.5. 

(2) and (3). 

39 ^ 122.5 = 0.3184, and 0.3184 X 100 = 31.84% of K. 

35.5 -4- 122.5 = 0.2898, and 0.2898 X 100 = 28.98% of Cl. 

48 122.5 = 0.3918, and 0.3918 X 100 = 39.18% of O. 

This process of finding the composition in per cent from 
the formula is often called calculating the percentage com¬ 
position of a compound. 

105. The simplest formula of a compound can be calcu¬ 
lated from its percentage composition. — The calculation of 
a formula from the percentage composition is simply the 
reverse of the process given in § 104, i.e., thus: — 

(1) Divide each per cent by the corresponding (single) 
atomic weight. 

(2) Reduce the quotients to whole numbers. 

(3) Write the formula. 

Let us take an example. The composition of sulfuric 
acid is hydrogen = 2.04 per cent, sulfur = 32.65 per cent, 
oxygen = 65.31 per cent. 

(1) Dividing the percentage of each element by the 
corresponding atomic weight, the quotients are 2.04, 1.02, 
and 4.08. 

(2) Reducing these quotients to whole numbers (by divid¬ 
ing by 1.02 in this case), the final quotients are 2, 1, 4. 

(3) These quotients represent the number of atomic 
weights of the respective elements in a molecule of this 
compound. And since atomic weights and atoms are repre¬ 
sented by symbols, a molecule of sulfuric acid contains 
2 atoms of hydrogen (2 H), 1 of sulfur (S), and 4 of oxygen 
(4 O). This means that the formula of sulfuric acid must 
be H2SO4. 

Formulas calculated by this method are often the simplest 
formulas of the corresponding compound. 


SYMBOLS AND FORMULAS 


83 


EXERCISES 

1. Learn the symbols of these elements: (a) Aluminum, calcium, 
carbon, chlorine, copper, (h) Hydrogen, iron, lead, magnesium, mer¬ 
cury, nitrogen, oxygen, (c) Phosphorus, potassium, silicon, silver, 
sodium, sulfur, zinc. 

2. What elements correspond to C, Cl, Ca, Cu, S, Si, Mg, Hg, H ? 

3. What do these mean? H, H 2 , 2 H, 2 O 2 , Cl, CI 2 , 3 Cl, 3 CI 2 , N 2 , 
K, 2 Ca, 3 Fe, S 2 , Cu, AL. 

• 4. What do these mean? H 2 O, 2 H 2 O, KNO 3 , 4 H 2 SO 4 , NaOH, 

3 Ca(OH) 2 , HNO 3 , BaCL . 2 H 2 O, 2 FeS, 3 CaCL. 

(Hint. H 2 O means 1 molecule of water containing 2 atoms of hydro¬ 
gen combined with 1 atom of oxygen.) 

6 . What is the formula of water, potassium chlorate, sulfuric acid, 
magnesium oxide, copper oxide, sodium hydroxide ? What is the molec¬ 
ular weight of each? 

PROBLEMS 

(Some of these problems may be used in subsequent assignments.) 

1. Calculate the molecular weight (or multimolecular weight) of 
the following compounds by finding the sum of the atomic weights: 

(а) magnesium oxide (MgO), (6) hydrogen peroxide (H 2 O 2 ), (c) zinc 
chloride (ZnCL), (d) 2 Cu(N 03 ) 2 , (e) 3 Al 2 (S 04 ) 3 , (/) FeS 04 .7 H 2 O. 

2 . Calculate the simplest formula of each compound from the indi¬ 
cated percentage composition: (a) Cl = 60.68, Na = 39.31; (6) S = 
23.52, Ca = 29.41, O = 47.05; (c) C = 40, H = 6.67, O = 53.33. 

3. As in Problem 2: (a) N = 26.17, H = 7.48, Cl = 66.35; 

(б) As = 75.8, O = 24.2; (c) N = 82.35, H = 17.65. 

4. As in Problem 2: (a) Si = 19.5, C = 66.62, H = 13.88; (6) Ca 

= 38.71, P = 20, O = 41.29; (c) H = 1, K = 39.06, C = 11.99, 

O = 47.95. 

6 . Calculate the formula of a compound 18 gm. of which contain 
8.4 gm. of iron and 9.6 gm. of sulfur. 

6 . As in Problem 6 : 0.84 gm. contain 0.587 gm. of iron and 0.253 
gm. of oxygen. 

7. Calculate the percentage composition of (a) hydrochloric acid, 
(6) hydrogen sulfide (H 2 S), (c) ammonia (NH 3 ), (d) hydrogen peroxide 
(H 2 O 2 ). 

8. As in Problem 7: (a) calcium oxide (CaO), (&) calcium carbon¬ 
ate (CaC 03 ), (c) calcium sulfate (CaS 04 ), (d) calcium fluoride (CaF 2 ). 

9. As in Problem 7: (a) cane sugar (C 12 H 22 OU) and (6) grape sugar 

(C6H,206). 

10. As in Problem 7: (a) sodium phosphate (Na 3 P 04 ), (&) disodium 
phosphate (Na 2 HP 04 ), (c) monosodium phosphate (NaH 2 P 04 ), (d) 
phosphoric acid (H 3 PO 4 ). 

11. Calculate the per cent of (a) copper and (b) water in crystallized 
copper sulfate (CUSO 4 .5 H 2 O). 


84 


A BRIEF COURSE IN CHEMISTRY 


SUPPLEMENTARY SECTIONS FROM PART II 

363. Facts, laws, and theories (see § 346). 

364. Law of the conservation of matter. 

366. Law of constant composition. 

366. The atomic theory. 

367. Atoms and molecules. 

368. Interpretation of chemical change by the atomic theory. 

369. How the atomic theory helps us explain fundamental laws. 

360. Atomic weights. 

361. Equivalent weights. 

362. The difference between equivalent and atomic weights. 


CHAPTER IX 


CHEMICAL REACTIONS — EQUATIONS — 
CALCULATIONS 

106. Chemical reactions. — A chemical change is called a 
chemical reaction, an interaction, or simply a reaction. 

107. What is a chemical equation? — We learned in 
§ 17 that a reaction can be represented in a condensed form 
by a verbal equation. Thus: — 

Iron + Sulfur = Iron Sulfide 

In §§ 14, 16 we learned that elements and compounds are 
represented by symbols and formulas. Therefore, we can 
use symbols and formulas in place of words. The above 
equation then becomes : — 

Fe + S = FeS 

108. How to read an equation. — The plus (+) sign may 
be read and or 'plus and the equality ( = ) sign form(s), 
give(s), yield{s), or equates). An arrow (->-) is sometimes 
used instead of the equality sign (=); both signs are read 
in the same way. Since equations are really expressions of 
equality between two total weights, the sign of equality is 
the proper sign to use. (However, see § 192.) 

Consider the equation: 

Zn + H2SO4 = H2 + ZnS04 

This equation may be read in several ways: — 

(1) Zinc and sulfuric acid form (or give) hydrogen and 
zinc sulfate. 

(2) Zinc plus sulfuric acid equal hydrogen plus zinc 
sulfate. 


85 


86 


A BRIEF COURSE IN CHEMISTRY 


(3) One atom of zinc and one molecule of sulfuric acid 
form one molecule of hydrogen and one molecule of zinc 
sulfate. 

(4) Anticipating a little (§ 112), 65 parts of zinc interact 
with 98 parts of sulfuric acid and yield 2 parts of hydrogen 
and 161 parts of zinc sulfate, all parts by weight. 

109. How to write an equation. — The simplest way is to 
substitute symbols and formulas for the names of substances. 
But the mere change of words to symbols and formulas would 
not always give a correct equation. So these steps should 
be followed in writing an equation : — 

(1) Find (in the Table of atomic weights or in the text (by 
the Index)) the correct symbol or formula of each substance 
involved in the reaction. 

(2) Write a preliminary equation by putting on the left 
(of the equality sign) the symbol or formula of each factor, 
i.e., the original, initial, or reacting substances, and on the 
right the symbol or formula of each product (^.e., the final 
substances). 

(3) Balance the preliminary equation (if necessary), i.e., 
increase the number of atoms, or molecules, or both, as 
necessary, until the number of atoms of each element (free 
or combined) is the same on each side of the equation (§ 110 ). 

(4) Check the final equation. (See § 110.) 

110. Writing equations. — Let us take several examples 
of the simplest way. 

1. Magnesium combines with oxygen to form magnesium 
oxide. 

(1) The symbols and formulas are Mg, O 2 , and MgO. 
The last is found (if not remembered) by looking up the 
name (magnesium oxide) in the Index and referring to the 
page where this substance is described. 

(2) The preliminary equation is : — 

Mg + O 2 = MgO 

(3) Inspection shows 2 atoms of oxygen (in O 2 ) are on the 
left and only I (in MgO) is on the right. Hence the pre¬ 
liminary equation must be balanced, i.e., we must increase 
the number of oxygen atoms on the right. We do this by 


CHEMICAL REACTIONS 


87 


prefixing the coefficient 2 to the formula MgO, not hy altering 
the formula MgO. This change increases the Mg to 2 (be¬ 
cause a coefficient multiplies a whole formula) and necessi¬ 
tates multiplying Mg on the left by 2. The balanced equa¬ 
tion then becomes: — 

2 Mg + O 2 = 2 MgO 

(4) Final checking shows the same number of atoms of 
each element on both sides of the equation, hence the equation 
is correct. 

2. Zinc and hydrochloric acid form hydrogen and zinc 
chloride. 

(1) and (2) The preliminary equation is: — 

Zn + HCl = H 2 + ZnCl 2 

(3) Inspection shows 2 atoms of chlorine (in ZnCh) are 
on the right and only 1 is on the left. Hence we must 
multiply HCl by 2, thereby providing 2 atoms of chlorine, 
and giving also 2 atoms of hydrogen. The balanced equa¬ 
tion becomes: — 

Zn + 2 HCl = H 2 + ZnCl 2 

(4) Checking shows this equation is correct. 

3. Potassium chlorate decomposes into oxygen and 
potassium chloride. 

(1) and (2) The preliminary equation is: — 

KCIO 3 = 02 + KCl 

(3) Inspection shows 3 atoms of oxygen (in KCIO 3 ) are 

on the left and only 2 (in O 2 ) are on the right. So we multi¬ 
ply KCIO 3 by 2 and O 2 by 3. The second preliminary equa- 
tion is: — ^ = 3 O 2 + KCl 

Inspecting again, it is obvious there are 2 K and 2 Cl on the 
left, but only 1 of each is on the right. Therefore, we bal¬ 
ance by multiplying KCl by 2, and obtain as a final equa- 

tion:— 2 KClOs = 3 O 2 + 2 KCl 

(4) The equation checks, and is, therefore, correct. 


88 


A BRIEF COURSE IN CHEMISTRY 


111. Some precautions to be observed in writing equations 
by the above method. — (a) Correct formulas must be used. 
If a formula is not remembered or is not given in the im¬ 
mediate text, do not guess. Look it up in the book by find¬ 
ing the name of the substance in the Index and consulting 
the proper page where, as a rule, the formula is given. 

(6) It must not be overlooked that the correct formulas of 
many elementary gases, such as oxygen, hydrogen, nitrogen, 
and chlorine, are O2, H 2 , N2, CI2 respectively (not O, H, etc.). 

(c) Only the substances that actually take part in the 
chemical change should be included in the equation. Thus, 
when magnesium is burned in air, the nitrogen of the air 
does not unite with the magnesium (to any extent). Hence 
nitrogen does not appear in the equation. Similarly in the 
equation for the preparation of hydrogen from zinc and 
hydrochloric acid, no water (H 2 O) appears as a factor be¬ 
cause the water (in the dilute acid), being only a solvent, does 
not participate in the reaction. And in the equation for the 
preparation of oxygen from potassium chlorate and man¬ 
ganese dioxide, the manganese dioxide is omitted because 
it acts only as a catalyst. 

(d) In balancing an equation these rules are helpful: — 

(1) Start with the formula containing the most atoms of 
one element. 

(2) Find the other formula (or symbol) containing this 
element, and increase the number of atoms of this element 
by prefixing a coefficient, not by altering the formula (or 
symbol) — see § 110 1 (3). 

(3) Increase again, if necessary. 

(4) Balance in the same way for the other elements. 

(5) Check up, so that finally the total number of atoms of 
each element is the same on both sides. 

Let us take an example. When phosphorus burns in 
oxygen, phosphorus pentoxide is formed. The formula of 
phosphorus pentoxide is P2O5. Hence the preliminary 
equation is:- P + 0, = P.Os 

By inspection we see that P 2 O 5 needs at least 5 atoms of 
oxygen. Clearly the only way to balance the equation for 


CHEMICAL REACTIONS 89 

oxygen is to multiply O 2 by 5 and P 2 O 5 by 2; this adjustment 
gives 10 atoms of oxygen on each side, thus: — 

P + 5 O 2 = 2 P 2 O 5 

But multiplying P 2 O 5 by 2 gives 4 atoms of P on the right, 
because a coefficient multiplies the whole formula. That is, 
2 P 2 O 6 means two molecules each containing 2 P and 5 O. 
Hence we balance for P by multiplying P on the left by 4. 
The final equation then becomes: — 

4 P + 5 O2 = 2 P2O5 

112. Calculations from equations. — Equations enable us 
to find the weights of the substances that can be obtained 
from (or are needed for) a given weight of a single substance. 
Knowing one weight, and the proportions by weight in 
which the whole reaction occurs, we can readily calculate 
all other weights involved. 

Consider the reaction between zinc and sulfuric acid. The 
complete equation is: — 

Zn + H 2 SO 4 = H 2 + ZnS 04 

65 2 + 32 + 64 2 65 + 32 + 64 

65 98 2 161 

The equation in this form is read: 65 parts by weight of 
zinc plus 98 parts by weight of sulfuric acid equal 2 parts by 
weight of hydrogen and 161 parts by weight of zinc sulfate. 
That is, these numbers are the relative weights of the dif¬ 
ferent substances involved in this reaction, viz., zinc and 
sulfuric acid in the ratio of 65 to 98 and hydrogen and zinc 
sulfate in the ratio of 2 to 161. 

In actual practice we need not start with exactly 65 gm. 
of zinc or with 98 gm. of sulfuric acid. We can use any 
convenient weights; but whatever weights we use, these 
two substances interact in the ratio of 65 to 98 and the 
quantity of either substance greater than the amount for 
the required ratio will be left unchanged. 

Suppose the problem to be solved is this: What weight of 
sulfuric acid is needed to interact completely with 45 gm. of 
zinc? We follow strictly these steps: — 


90 


A BRIEF COURSE IN CHEMISTRY 


(1) Write the chemical equation. Thus: — 

Zn + H 2 SO 4 = H 2 + ZnS 04 

(2) Write the relative weight equation by placing under 
each symbol and formula the weight for which the complete 
symbol or formula stands (using the approximate atomic 
weights in the Table on the inside back cover). Thus: — 

Zn + H 2 SO 4 = H 2 + ZnS 04 
65 2 + 32 + 64 2 65 + 32 + 64 

65 98 2 161 

(3) Write the reacting weight equation by placing the 
weight given in the problem above the proper symbol or for¬ 
mula (45 gm. above Zn in this case) and x gm. above the for¬ 
mula of the required substance (sulfuric acid here). Thus: — 

45 gm. X gm. 

Zn + H 2 SO 4 = H 2 + ZnS 04 

65 98 2 161 

(4) State in a proportion the four quantities involved. 
In this proportion the equation weights (below, in the equa¬ 
tion) are the first and second terms, while the corresponding 
known and required weights (above, in the equation) are 
the other two terms. Thus: — 

65 : 98:: 45: X 

(5) Solve the proportion for x, remembering that the 
product of the means (the two inner terms) equals the prod¬ 
uct of the extremes (the two outer terms). Thus: — 

QQ V AK 

98 X 45 = 65 X x. a: = , or 67.8. Ans. 67.8 gm. 

65 

Instead of steps (3) and (4) an alternative method may be 
used. Thus, if 65 gm. of zinc require 98 gm. of sulfuric 
acid, 1 gm. of zinc will require 98 65, or 1.507, and 45 

gm. would require 45 X 1.507, or 67.8 gm. 

Problems involving volume are readily solved. Let us 
take two examples. 



CHEMICAL REACTIONS 


91 


1. We start with 15 gm. of potassium chlorate and wish 
to know the volume of oxygen which nan be obtained. 
We find the weight, first, by proceeding as above. Thus: — 

(1) Chemical equation: — 

2 KClOs = 3 O 2 -h 2 KCl 

(2) Relative weight equation : — 

2 KCIO 3 = 3 O 2 + 2 KCl 

2(39 + 35.5 -f 48) 3(32) 2(39 + 35.5) 

2(122.5) 96 2(74.5) 

245 96 149 

(3) Reacting weight equation : — 

15 iC 

2 KCIO 3 = 3 O 2 + 2 KCl 
245 96 149 

(4) Proportion: — 

245 : 96 :: 15 : X 

(5) Solution for weight: - 
96 X 15 = 245 X x. 

(6) Solution for volume: 

Since 1 liter of oxygen weighs 1.429 gm., 

5.87 1.429 = 4.1. Ans. 4.11. 

2. We wish to know the weight of calcium carbonate 
needed to produce 200 1. of carbon dioxide. First, we find 
the weight of 200 1. of carbon dioxide. Since 1 1. weighs 
1.98 gm., the weight of 200 liters is 200 X 1.98, or 396 gm. 
To find the weight of calcium carbonate needed for 396 gm. 
of carbon dioxide, we proceed as follows: — 

(1) CaC 03 + 2HC1 = CO 2 + CaCl 2 + H 2 O 

(2) CaC 03 + 2 HCl = CO 2 -f- CaCh + H 2 O 

40 + 12-1- 48 2(1 + 35.5) 12 + 32 40 + 71 2 + 16 

100 73 44 111 18 


, X — 


_ 96 X 15 


245 


or 5.87 



A BRIEF COURSE IN CHEMISTRY 


92 


(3) 

X 


396 




CaCOs 

+ 2 HCl 

= CO 2 

+ CaCl 2 

+ H 2 O 


100 

73 

44 

111 

18 

(4) 


0 

0 

: X : 396 



(5) 

44: X X = 

100 X 396. 

100 
X = — 

X 396 

-, or 

44 

900 Ans. 


Hence in solving problems involving volume we simply 
find (1) the volume corresponding to the calculated weight 
or (2) the weight equal to a given volume, and then solve for 
the required weight. 


EXERCISES 

1 . Write equations for the following reactions: (a) Calcium and 
hydrochloric acid form calcium chloride and hydrogen. (6) Potassium 
sulfate and barium chloride form barium sulfate and potassium chloride, 
(c) Calcium carbonate and hydrochloric acid form calcium chloride, 
water, and carbon dioxide. 

2 . As in Exercise 1 : (a) Calcium oxide and carbon dioxide form 
calcium carbonate. (6) Chlorine and aluminum form aluminum tri¬ 
chloride. (c) Carbon and lead oxide (PbO) form lead and carbon mon¬ 
oxide. 

3 . Balance these equations: (a) BaCh + H 2 SO 4 = BaS 04 + HCl; 
(b) Pb(N 03)2 + H 2 S = PbS + HNO 3 ; (c) AICI 3 + NH 4 OH = A 1 ( 0 H )3 
+ NH 4 CI; (d) NaOH + CO 2 = Na 2 C 03 + H 2 O. 

4 . Balance these equations : (a) HCl + ZnO = ZnCb + H 2 O ; 

(b) H2SO4 + NaN 03 = Na 2 S 04 + HNO3; (c) SO2 + 02 = SO3. 


PROBLEMS 

(Some of these problems may be used in subsequent assignments.) 

1 . How many grams of oxygen can be prepared from (a) 45 gm. of 
mercuric oxide, (b) 1 kg. of potassium chlorate, (c) 1000 gm. of water? 

2 . As in Problem 1 , from (a) 750 gm. of lead dioxide (Pb02), (b) 2200 
gm. of barium dioxide (Ba 02 ) ? 

3 . Hydrogen is prepared from sulfuric acid and 40 gm. of zinc. 
Calculate the weights of the products of the reaction. 

4 . If a balloon holds 150 kg. of hydrogen, how much (a) zinc and 
(6) sulfuric acid are needed to generate the gas? 

6. What volume of oxygen (at 0° C. and 760 mm.) can be obtained 
from 10 gm. of potassium chlorate? 



CHEMICAL REACTIONS 93 

6. If 10 gm. of pure carbon are burned in air, what weights of other 
substances are involved? 

7 . One gm, of copper is heated intensely in air, and the product is 
reduced by a gas. Calculate (a) the weights of the other substances in¬ 
volved in the two reactions, and (5) the volume of the gas in the second 
reaction. 

8. How many grams of potassium chlorate are needed to prepare 
(a) 100 gm. of oxygen and (b) 100 1. (at 0° C. and 760 mm.) ? 

9 . Calculate the weights needed in the following reactions: 

(а) water and 100 mg. of sodium, (b) calcium and 100 mg. of water, 
(c) sodium hydroxide and 25 gm. of aluminum. 

10. What weight of carbon dioxide is formed by burning a ton of 
coal which is 90 per cent carbon ? 

11. Suppose 85 gm. of water are decomposed. What (a) weights and 

(б) volumes of gases are produced? 

12. Sixty grams of mercuric oxide are decomposed. What volume 
of oxygen (at 0° C. and 760 mm.) is produced? 

13. How much water can be obtained from (a) 34 gm. of crys¬ 
tallized zinc suKate (ZnS 04 .7 H 2 O), (6) 1000 kg, of crystallized cal¬ 
cium sulfate (CaS 04.2 H 2 O), (c) 1000 gm. of washing soda crystals 

(NaaCOa. IOH 2 O)? 

14. Write the equation for the interaction of barium nitrate and 
sodium sulfate. If 170 gm. of barium nitrate are used, calculate the 
weights of the other compounds involved. 

16. Ammonia and hydrogen chloride form solid ammonium chloride. 
Write the equation for this reaction. If 210 gm. of ammonia are used, 
calculate (a) the weights of the other compounds involved and (6) the 
volumes of ammonia and hydrogen chloride needed. 

16. The oxygen is liberated from 10 gm. of potassium chlorate, and 
10 gm. of sulfur are burned in the gas. How much sulfur, if any, is left? 


CHAPTER X 


VALENCE 

113. What is valence ? — Atoms and radicals form 
molecules. This is merely another way of saying atoms 
and radicals have a certain chemical property, viz.^ a capacity 
to combine and displace. The number which expresses 
this combining or replacing capacity of an atom or a radical 
is called its valence. Usually we speak of the valence of an 
element or a radical, e.g., the valence of hydrogen is 1 and 
of SO4 is 2, though we mean the combined element or radical. 
Free or uncombined elements have the valence of zero. 

114. Hydrogen is the standard for reckoning valence. — 
No atom has a combining or replacing capacity lower than 
a hydrogen atom. Hence the hydrogen atom is the standard 
and its valence is 1. This means that 1 hydrogen atom holds 
finally in combination only 1 atom or 1 radical. Thus, in 
HCl and HNO3 1 atom of hydrogen is combined with 1 Cl 
atom and 1 NO3 radical. 

115. Classification of elements and radicals by their 
valence. — Several elements and radicals have the same 
combining and replacing capacity as hydrogen, i.e., their 
atoms and the radicals combine with or displace 1 hydrogen 
atom. The valence of these elements and radicals is 1, and, 
like hydrogen, they are called univalent elements and radicals, 
e.g., chlorine, sodium, and the nitrate (NO3) and hydroxyl 
(OH) radicals. (See Tables in § 117.) 

Many elements and radicals have the combining and 
replacing capacity of 2 hydrogen atoms, i.e., their valence 
is 2 and they are called bivalent elements and radicals, e.g., 
oxygen, calcium, and the sulfate radical (SO4). Similarly, 

94 


VALENCE 


95 


there are trivalent elements and radicals, whose valence is 3, 
e.g., aluminum and the phosphate radical (PO 4 ); quadriva¬ 
lent, whose valence is 4, e.g., carbon; quinquivalent, whose 
valence is 5, e.g., phosphorus, and so on up to 8 — the 
limit. 

Atoms and radicals which have the valence 1 are some¬ 
times called monads; similarly, we have dyads, triads, 
tetrads, 'pentads, etc. 

116. Positive and negative valence. — There are two 
kinds of valence — positive and negative which are desig¬ 
nated by the signs -|- and — respectively. This division 
corresponds in general to the metallic or non-metallic char¬ 
acter of the element. Thus, metals, e.g., sodium and cal¬ 
cium, have a positive valence. Whereas, non-metals, e.g., 
oxygen and chlorine, have a negative valence. The radical 
ammonium is positive, whereas the other radicals are 
negative. 

The division according to metals and non-metals is helpful. 
A better basis for division is the electrical behavior, e.g., in 
electrolysis, which will be studied in §§ 180-182. There we 
shall learn that if an element or radical goes to the cathode 
(negative electrode of the electrolytic apparatus), the valence 
is positive, but if it goes to the anode (positive electrode), the 
valence is negative. 

On the electrical basis, for example, the valence of com¬ 
bined hydrogen is + 1, oxygen is — 2, calcium is + 2, and 
the S 04 -radical is — 2. These valences are sometimes 
written H+b OS Ca+ 2 , BOS. 

The usual valence of the common elements and radicals 
in compounds is shown in the accompanying table. We 
should always remember that the valence of a free (i.e., un¬ 
combined) element is zero, e.g., H°, 2 CP, N 2 °. Hence, the 
valence in the table means that these symbols and formulas 
have the designated valence in compounds, e.g., free hydrogen 
is but combined hydrogen is H+b and (by analogy) free 
OH is OH° but OH in NaOH is OH“L 

117. Tables of valence. — The usual valence of the com¬ 
mon elements and radicals is shown in Table III, A and B. 
These tables are very useful in writing formulas (§ 118). 


96 


A BRIEF COURSE IN CHEMISTRY 


TABLE III. — A. Positive Valence of Elements and 
Radicals 


Name 

Symbol 

Positive 

Valence 

Name 

Symbol 

Positive 

Valence 

Aluminum 

A1 

+ 3 

Lead 

Pb 

+ 2 

Ammonium 

NH 4 

+ 1 

Magnesium 

Mg 

+ 2 

Barium 

Ba 

+ 2 

Mercury (ous) 

Hg 

+ 1 

Calcium 

Ca 

+ 2 

Mercury (ic) 

Hg 

+ 2 

Carbon 

C 

+ 4 

Phosphorus (ic) 

P 

+ 5 

Copper (ic) 

Cu 

+ 2 

Potassium 

K 

+ 1 

Gold (ic) 

Au 

+ 3 

Silicon 

Si 

+ 4 

Hydrogen 

H 

+ 1 

Silver 

Ag 

+ 1 

Iron (ous) 

Fe 

+ 2 

Sodium 

Na 

+ 1 

Iron (ic) 

Fe 

+ 3 

Zinc 

Zn 

+ 2 


TABLE III. — B. Negative Valence of Elements and 
Radicals 


Name 

Symbol 

OR 

Formula 

Valence 

Name 

Symbol 

OR 

Formula 

Valence 

Bromine 

Br 

- 1 

Iodine 

I 

- 1 

Carbonate 

CO 3 

- 2 

Nitrate 

NO, 

- 1 

Chlorate 

CIO, 

- 1 

Oxygen 

0 

- 2 

Chlorine 

Cl 

- 1 

Permanganate 

Mn04 

- 1 

Dichromate 

Cr207 

- 2 

Phosphate 

PO 4 

- 3 

Ferrocyanide 

Fe (CN)6 

- 4 

Sulfate 

SO 4 

- 2 

Hydroxide 

OH 

- 1 

Sulfate (acid) 

HSO 4 

- 1 


Some elements have more than one valence because they 
form different kinds of compounds. Thus, the valence of 
iron is + 2 in ferrous compounds and + 3 in ferric com¬ 
pounds, e.g., ferrous sulfate, Fe+ 2 S 04 ~^ and ferric chloride 
Fe+^Cl 3 “^“^“h Similarly, sulfur is S^^ in sulfides, S"^ in 
sulfur dioxide, and S+® in sulfur trioxide. In general, the 
valence in Tables A and B is the numerical value in com¬ 
mon compounds, particularly compounds of two elements or 



























VALENCE 97 

of an element and a radical, e.g., Na+iCl“S Ca+ 2 S 04 -^ 
Cu+2(N03)2-1-^ 

118. Writing formulas by positive and negative valence. 

— The formulas of many compounds can be written by 
utilizing two rules: — 

1. Atoms and radicals of the opposite kind of valence 
combine, i.e., positive with negative. 

2. The positive and the negative parts must be numerically 
equal, i.e., the valence must balance and the algebraic sum 
must be 0. 

Let us take some examples. 

(а) What is the formula of calcium sulfate? From Table 
A calcium is + 2, and from Table B the sulfate radical is 

— 2. Hence the formula is Ca+ 2 S 04 ~^ or CaS 04 . 

(б) What is the formula of aluminum oxide? From 
Table A aluminum is + 3 and from Table B oxygen is 

— 2. To make the parts balance, we need 2X(H-3) = +6 
and 3 X(— 2)= — 6, or 2 A1 and 3 0; hence we write 
AI 2 + 3 + 3 O 3 + 2 + 2+2 or AI 2 O 3 . 

(c) How are formulas of salts derived? Formulas of 
salts are frequently needed. These can be readily written 
by learning the formulas of the common acids and sub¬ 
stituting the correct number of atoms for the hydrogen. 
Thus, the formula of sulfuric acid is H 2 SO 4 . Therefore, all 
sulfates will have one or more S 04 -radicals as one part of the 
formula and one or more atoms of a metal for the other part. 
If the metal has the valence of -f- 1, the formula will contain 
two atoms of the metal, e.g., Na 2 S 04 ; if the valence of the 
metal is + 2 , there will be only one atom, e.g., CUSO 4 . 
Whereas if the valence is + 3, there will be two atoms of the 
metal and three S 04 -radicals, e.g., AI 2 ( 804 ) 3 . Similarly, 
all nitrates have NO3 (one or more), all chlorides have Cl 
(one or more), all phosphates have PO 4 (one or more). 

119. Representation of valence. — The valence of ele¬ 
ments and radicals may be represented in various ways. 
One has already been given (exponent with proper sign), 
e.g., H+h 804“^. Sometimes short lines are used, e.g., 
H—, 0=, —O—, Al=, etc. If lines are used to represent 
valence in compounds, a single line answers for two elements. 


98 


A BRIEF COURSE IN CHEMISTRY 


Thus, the formula of water is written H—O—H. Similarly, 

the formulas of ammonia, calcium hydroxide, and nitric acid 
OH .0—H ^ 

or Ca<^ 

OH \0—H 

Such formulas are called structural or graphic formulas, 

for they show the probable arrangement of the atoms in the 
molecules. It must not be forgotten that structural formulas 
are merely representations; the lines are intended to indicate 
the numerical value of the valence and not the strength of 
the combination of the atoms. 

EXERCISES 

(Some of these Exercises may be used in subsequent assignments.) 

1. What is the valence of; (a) H, O, C, Cl; ( 6 ) Ag, K, Na, 
NH 4 ; (c) NO3, OH; (d) Ba, Ca, Cu, Fe (ous), Mg, Pb, Zn; (e) CO3, 
SO4, S (ide); (/) Al, Fe (ic), PO4? 

2 . State the fundamental rules for writing formulas from valence. 

3 . Write the formula of the chloride of potassium, sodium, silver, 
copper, mercury (ous), mercury (ic), iron (ous), iron (ic), zinc, calcium, 
barium, magnesium, aluminum, ammonium, lead. 

4. Write the formula of the sulfate of K, Na, Ag, Cu, Fe (ous), Zn, 
Pb, Ca, Ba, Mg, Al, NH4. 

6 . Write the formula of the following compounds and indicate 
the valence of each element or radical: ammonium chloride, potassium 
carbonate, barium phosphate, zinc iodide, ammonium dichromate, 
silver sulfate, magnesium oxide, sodium dichromate, aluminum chloride, 
ferrous bromide, calcium phosphate, mercurous nitrate. 

6 . Write the formulas of the following (as in Exercise 6 ): ferrous 
carbonate, aluminum phosphate, calcium iodide, sodium permanganate, 
phosphoric acid, carbonic acid. 

7. Write the formula of the nitrate of calcium, silver, iron (ic), 
mercury (ous), barium, magnesium, aluminum. 

8 . Write the formula of the oxide of Al, Ba, Ca, Na, Ag, K, Zn, Mg. 

9. Write the formula of (a) the carbonate of Cu, Ba, Ag ; ( 6 ) the 
iodide of Hg (ous and ic), Pb, Ag, Na, Fe (ous and ic), Al; (c) the chlo¬ 
rate of Na, Ba, Ag. 

10. Write the formula of (a) the dichromate of Pb, Na, Al, Ag; 
(b) the phosphate of Na, Ca, Ag; (c) the nitrate of Pb, Ca, Cu. 

PROBLEMS 

See Problems at the end of Chapters VIII and IX. 


, H—0—N: 




O 


are: 


N^H, Ca< 
\H 


CHAPTER XI 


CHLORINE — HYDROGEN CHLORIDE 
— HYDROCHLORIC ACID 

120. Occurrence of chlorine. — Free chlorine is never 
found in nature because it is such an active chemical element. 
But some of its compounds, especially chlorides, are widely 
distributed. The most abundant compound is sodium chlo¬ 
ride (NaCl) or common salt. 

Thick beds of compounds of 
chlorine with potassium, mag¬ 
nesium, and calcium occur in 
the Stassfurt region in Germany 
and the Alsace-Lorraine district 
in France; some beds are almost 
entirely potassium chloride. 

Ocean water and the water of 
such lakes as the Dead Sea and 
the Great Salt Lake contain 
a large proportion of sodium 
chloride. 

121. Manufacture of chlorine. 

— Chlorine is manufactured by 
the electrolysis of a solution of 
sodium chloride (Figs. 49, 50). 

When an electric current is 
passed through a solution of sodium chloride, chlorine is 
liberated in one compartment of the apparatus and sodium 
hydroxide is formed in the other. The chlorine gas is 
usually conducted off through a pipe {A in Fig. 49) and 
compressed to a liquid in strong metal tanks (Fig. 52). 

99 



facturing chlorine by the elec¬ 
trolysis of a solution of sodium 
chloride. The chlorine escapes 
through the pipe A. 




















100 


A BRIEF COURSE IN CHEMISTRY 


The sodium hydroxide dissolves and the solution is drawn 
off at intervals. This process is more fully described in 

§ 285. 

The manufacture of chlorine can be readily 
demonstrated in the laboratory. The apparatus 
is shown in Fig. 50. A solution of sodium chlo¬ 
ride is put in the battery jar A; a little litmus 
solution is added and then enough dilute hydro¬ 
chloric acid to color the solution a distinct red. 
A block B divides the jar into two compart¬ 
ments C and D, and the two pieces of electric 
light carbon serve as electrodes E and F. Soon 
after the current (from a storage battery or four 
or more dry cells) is turned on, the solution is 
bleached by the liberated chlorine in one com¬ 
partment and turned blue by the sodium hy¬ 
droxide in the other. The chlorine can also be 
detected by its choking odor. 

of chlorine in the laboratory. — Chlorine 
is prepared in the laboratory by heating concentrated hydro¬ 
chloric acid with an oxidizing agent, 
such as potassium permanganate 
(KMn 04 ) or manganese dioxide 
(Mn02). The equation for the 
reaction with manganese dioxide 
is: — 

4 HCl + Mn02 

Hydrochloric Acid Manganese Dioxide 

CU + MnCU + 2 H 2 O 

Chlorine Manganese Dichloride Water 

If only a small quantity of chlorine is 
wanted, enough crystallized potassium 
permanganate is put into a warm bottle 
to form a thin layer, 5 cc. of concentrated 
hydrochloric acid is added, and the bottle 
is covered with a piece of filter paper. 

Chlorine (gas) soon fills the bottle. If a 
larger quantity is needed, the apparatus 
shown in Fig. 51 may be used. 

Manganese dioxide is put into the flask 
A and concentrated hydrochloric acid is 
introduced through the dropping tube B. 



Fig. 51. —Apparatus for pre¬ 
paring chlorine in the labo¬ 
ratory from hydrochloric 
acid and manganese dioxide. 

By heating the flask gently, 



Fig. 50. — Laboratory 
apparatus for the 
preparation of chlo¬ 
rine by the electroly¬ 
sis of a solution of 
sodium chloride. 

122. Preparation 


































CHLORINE — HYDROGEN CHLORIDE 101 

chlorine passes to the bottom of the bottle G, rises slowly, and dis¬ 
places the air. 


123. Physical properties of chlorine. — Chlorine is a 
greenish yellow gas. It has a stifling odor. If breathed, it 
irritates the lining of the nose and throat; a large quantity 
produces suffocation, and would ultimately cause death. 
Chlorine was the first poison gas used in the World War. 
When chlorine is used in manufacturing or purifying plants, 
a gas mask is usually worn 
(Fig. 52). 

Chlorine is a heavy gas — 
about 2.5 times heavier than 
air. A liter at 0° C. and 
760 mm. weighs 3.22 gm. 

Chlorine can be readily 
liquefied. Liquid chlorine is 
sold in strong metal cylinders 
(Fig. 52 and also Fig. 53). 

Chlorine dissolves in water. 

The solution is yellowish, and 
smells like chlorine. Chlorine 
water, as the solution is called, 
is unstable. If the solution is 
placed in the sunlight, oxygen 
is slowly liberated and can be 
collected in a suitable appa¬ 
ratus (Fig. 45, § 96). 

The oxygen does not come 
directly from the water but from a compound of chlorine 
called hypochlorous acid (HCIO), which is formed in small 
quantities in the solution; some hydrochloric acid is also 
formed. The reaction may be represented thus : — 

CI 2 + H 2 O = HCIO + HCl 

Chlorine Water Hypochlorous Acid Hydrochloric Acid 

The hypochlorous acid is unstable and decomposes, thus: — 
2 HCIO = 02 + 2 HCl 

Hypochlorous Acid Oxygen Hydrochloric Acid 







102 


A BRIEF COURSE IN CHEMISTRY 


The reactions continue until all the chlorine is used up; 
so the equation for the completed change is: — 

2 H 2 O + 2 CI 2 = 4HC1 + O 2 

Water Chlorine Hydrochloric Acid Oxygen 

124. Chemical properties of chlorine. — Chlorine com¬ 
bines vigorously with many elements and interacts with 
many compounds. And just as oxygen — another active 
element — forms oxides, so chlorine forms chlorides. 



Fig. 53. — Interior of a plant for liquefying and storing chlorine. Cylinders 
full of chlorine are in the foreground. 


The uniting of chlorine with elements is easily shown. 
Thus, if sodium, iron, copper, or other metals are slightly 
warmed and put into chlorine, they unite with the chlorine 
at once; the sodium produces a dazzling light, and the 
copper and iron glow and emit dense fumes. These chemical 
changes illustrate the broad use of the term combustion. No 
oxygen is involved. Chlorine and the metal unite, and the 
chemical change is attended with heat and light. The 



















CHLORINE — HYDROGEN CHLORIDE 


103 


compound formed in each case is a chloride, i.e., a compound 
of chlorine and one other element. Thus, sodium and iron 
form sodium chloride (NaCl) and iron chloride (FeCls) 
respectively. 

Chlorine combines readily with hydrogen. A jet of burn¬ 
ing hydrogen when lowered into chlorine continues to burn, 
just as it does in oxygen. The product is a colorless gas 
called hydrogen chloride, which becomes a white cloud in 
moist air or when the breath (containing moisture) is blown 
across the mouth of the vessel (§ 131, second paragraph). 

The activity of chlorine toward hydrogen is so great that 
chlorine withdraws hydrogen chemically from many com¬ 
pounds, e.g., turpentine (CioHie). Thus, when cotton 
saturated with hot turpentine is put into chlorine, white 
fumes, due to the formation of hydrogen chloride, appear 
at once. Soon the chemical change is so vigorous that a 
flame is suddenly produced and the white fumes of hydrogen 
chloride are obscured by a dense cloud of black smoke, 
which consists of fine particles of carbon (left over from the 
decomposed turpentine). 

125. Chlorine compounds are used for bleaching. — 

Moist chlorine (essentially hypochlorous acid — § 126) and 
some chlorine compounds change many colored substances 
into colorless or pale ones. Thus, for a century or more 
cloth made of cotton or linen, naturally colored a faint 
yellow by impure substances, has been bleached (Le., whit¬ 
ened) by exposure to a moist compound of chlorine called 
bleaching powder. 

For certain industries, e.g., bleaching cotton cloth, the 
best bleaching agent is bleaching powder (CaOCh or pref¬ 
erably CaCl(OCl)) or, as it is often called, chloride of lime. 
It is a yellowish white substance which smells like chlorine, 
though the smell is really due to chlorine compounds, which 
are slowly liberated from the bleaching powder by the moist 
carbon dioxide in the air. 

In using bleaching powder, rather dilute sulfuric acid is 
mixed with a weak solution (or suspension) of the powder. 
Hypochlorous acid (HCIO) is liberated slowly. This un¬ 
stable acid yields oxygen, which does the bleaching. 


104 


A BRIEF COURSE IN CHEMISTRY 


Another compound of chlorine used in bleaching is sodium 
hypochlorite (NaClO). This compound is prepared by 
passing an electric current through sodium chloride solution 
and allowing the products to react, thus: — 

2NaOH + CI 2 = NaClO + NaCl + H 2 O 

Sodium Chlorine Sodium Sodium Water 

Hydroxi4e Hypochlorite Chloride 

Sodium hypochlorite is unstable, and a cold dilute solution 
contains hypochlorous acid, thus : — 

NaClO + H 2 O = HCIO + NaOH 

Sodium Water Hypochlorous Sodium 

Hsqjochlorite Acid Hydroxide 

The hypochlorous acid, as just stated, decomposes and 
furnishes the oxygen. Sodium hypochlorite solution is used 
in bleaching wood pulp and for whitening linen in laundries. 

126. The process of bleaching. — The fibers of cotton 
and linen and also of the pulp obtained from wood consist 



Fig. 54. — Diagram of the process of bleaching cotton cloth. 


mainly of cellulose, which is not appreciably affected by 
dilute hypochlorous acid solutions. But the natural color¬ 
ing matter in these fibers is slowly oxidized to colorless 
compounds. Artificial coloring matter, e.g., dye, is usually 
decomposed (and decolorized) by hypochlorous acid. Some 
colored fabrics, e.g., those dyed with a mordant (§ 610), 
are not whitened completely, but only faded, by bleaching. 
Some dyes, e.g., those used for postage stamps, resist the 
bleaching action of chlorine compounds. Substances made 
up largely of carbon cannot be bleached, e.g., printer’s ink, 
“ lead ” of lead pencils, many canceling inks, and tarry 
mixtures from coal and petroleum. 

A diagram of the process of bleaching cotton cloth is shown in Fig. 54. 
The pieces are sewed together end to end in long strips and drawn by 
machinery from the roll A successively through vats containing bleach- 



























CHLORINE — HYDROGEN CHLORIDE 


105 


ing powder solution B, weak acid C, and water D. At some point 
toward the end of the process the cloth passes through a vat containing 
acid sodium sulfite solution (or similar mixture) E, called the antichlor, 
to remove traces of hypochlorous acid. After thorough washing, the 
cloth is dried and ironed by passing over hot cylinders F, G, and finally 
wound on the roll H. 

127. Uses of chlorine — Besides the use of chlorine in 
the manufacture of bleaching powder and bleaching mixtures, 
large quantities are made into useful compounds of chlorine. 
One is carbon tetrachloride (CCI4), which is used in “ py¬ 
rene fire extinguishers (Fig. 55) because it 
readily forms a heavy non-combustible vapor 
which smothers the fire, especially burning 
gasolene. Carbon tetrachloride is also used 
as a solvent for extracting greases. The non¬ 
combustible cleaning mixture called ‘‘ carbona ” 
contains carbon tetrachloride. Considerable 
chlorine is used to purify drinking water, and 
its use is increasing because it is a convenient 
and effective way to destroy bacteria (§ 81, 

Fig. 37). 

128. Chlorineinthe World War. —Chlorine and its Fig. 55 . — Fire 
compounds played a significant part in the World War. contarnliT^ 
Chlorine and phosgene (COCI2) were the first poison carbon tet- 
substances used. Tanks of liquid chlorine, or chlorine rachloride. 
and phosgene, were opened, the gas rushed out, and 

under favorable conditions was blown along the surface as a great cloud. 
Because it is heavier than air it sank into the trenches and immediately 
drove out the men. 

The cloud method was soon abandoned for the shell method, i.e., 
gases were replaced by readily volatilized substances, over haK being 
chlorine compounds. In fact chlorine was utilized, directly or indirectly, 
in preparing about 95 per cent of the toxic substances used to rout out 
the enemy. Some substances, e.g., chlorpicrin and benzyl chloride, 
caused a copious flow of tears (“ tear gases ”). Others, e.g., diphenyl- 
chlor-arsine, produced excessive sneezing (“ sneeze gases ”). Still 
others, e.g., mustard gas or di-chlor-ethyl-sulfide, irritated the skin and 
led to serious results. 

The effects of these substances were prevented or lessened by gas 
masks. The masks fitted the face closely, necessitating breathing 
through a canister. The canister contained a suitable absorbent, usually 
a dense variety of charcoal, prepared from cocoanut shells and fruit pits, 














106 


A BRIEF COURSE IN CHEMISTRY 


together with other substances, especially potassium permanganate 
and soda-lime (a special mixture of lime and sodium hydroxide). The 
soda-lime reacted with chlorine and chlorine compounds. 

Sihcon tetrachloride (SiCh) mixed with ammonia and water was used 
as a smoke screen, because these three reacted and formed a dense 
white cloud consisting of fine particles of ammonium chloride (NH4CI) 
and silicic acid (H 4 Si 04 ). This smoke settled down slowly and screened 
a vessel or the land (Fig. 56). 



Fig, 56. — A smoke screen laid by a warship off Cape Hatteras. 


HYDROGEN CHLORIDE AND HYDROCHLORIC ACID 

129. Hydrogen chloride and hydrochloric acid. — Hydro¬ 
gen chloride is a gas, which is very soluble in water. Hy¬ 
drochloric acid is the common name of a water solution of 
hydrogen chloride. This solution is known commercially as 
muriatic acid (from the Latin word muria, meaning brine), 
but it is more properly called hydrochloric acid. Hydrogen 
chloride is often called hydrochloric acid gas. 

130. Preparation of hydrogen chloride. — This gas is 
prepared in the laboratory by heating sodium chloride with 
sulfuric acid. If the mixture is heated gently, the chemical 
change is represented thus : — 

NaCl + H 2 SO 4 = HCl + NaHS 04 

Sodium Sulfuric Hydrogen Acid Sodium 

Chloride Acid Chloride Sulfate 




CHLORINE — HYDROGEN CHLORIDE 


107 


But at a high temperature the equation is: — 

2 NaCl + H2SO4 = 2 HCl + NasSOi 

Sodium Sulfate 

The gas is collected by passing it to the bottom of a bottle. 
The solution is prepared by absorbing the gas in water. 

Hydrochloric acid 
is manufactured in 
enormous quantities 
by a method essen¬ 
tially like that used 
in the laboratory. 

The mixture of salt 
and sulfuric acid is heated 
in a cast iron retort A 
(Fig. 57) by the furnace 
B to a moderate temper¬ 
ature ; as soon as the mass becomes pasty it is raked out upon the flat 
heater A' and heated to a high temperature by the furnace B'. The 
hydrogen chloride passes up through C and C' into an absorbing system 
(Fig. 58). It is cooled in the tower D, washed, cooled, and absorbed in 
the tower E, absorbed in the stoneware jugs FFF (through which weak 
acid from the absorbing tower flows), and finally absorbed in the tower 
G. The final portion is caught in the jug H. The concentrated solu¬ 
tion is drawn off at /. 


-S 

cH 


Fia. 68. — Sketch of the absorbing system used for the manufacture of 
hydrochloric acid. 

131. Physical properties of hydrogen chloride. — Hy¬ 
drogen chloride is a colorless gas. It has a choking, sharp 
taste, and irritates the lining of the nose and throat. It is 




Fig. 57. — Sketch of the furnace used for the 
manufacture of hydrochloric acid. 























































108 


A BRIEF COURSE IN CHEMISTRY 


about 1.25 times heavier than air. A liter at 0° C. and 
760 mm. weighs 1.64 gm. 

The solubility of hydrogen chloride in water is one of its 
most striking properties. When it escapes into moist air, 
it forms white fumes which are really minute drops of a 
solution of the gas in the moisture of the air. 


The solubility of hydrogen chloride in water can be shown by a simple 
experiment (Fig. 59). The flask A is filled with hydrogen chloride. 

The medicine dropper B is partly filled with 
water, the stopper with its tubes is inserted, 
and the flask is then arranged as shown in the 
figure. By pinching the bulb of the dropper, a 
few drops of water are forced into the flask. 
This small quantity of water dissolves so much 
gas that a partial vacuum is formed in the flask; 
pressure within the flask is reduced so much 
that the atmospheric pressure forces water from 
the jar C up the tube D and through the small 
opening into the flask. 


132. Chemical properties of hydro¬ 
gen chloride. — Hydrogen chloride does 
not burn nor support combustion. It 
is a very stable compound and can be 
heated to about 1800° C. before it 
begins to decompose. The moist gas 
unites readily with certain substances, 
e.g., ammonia; in this case dense white 
clouds of ammonium chloride (NH4CI) 
This reaction is used as a test for hydrogen 



Fig. 59. — Experiment 
to illustrate the 
marked solubility of 
hydrogen chloride. 


are formed 

chloride. 

133. Hydrochloric acid. — As previously stated, the 
water solution of hydrogen chloride is the common sub¬ 
stance hydrochloric acid, which is sometimes called muriatic 
acid. The commercial concentrated acid contains about 
35 per cent (by weight) of the actual substance HCl. Its 
specific gravity is about 1.2. The concentrated acid forms 
white fuines in the air, especially on a moist day, owing to the 
escape of hydrogen chloride (§ 131, second paragraph). 

Hydrochloric acid is a typical acid. It has a sour taste, 
reddens blue litmus, conducts electricity, and reacts with 




















CHLORINE — HYDROGEN CHLORIDE 


109 


most metals, forming hydrogen and salts. The salts are 
chlorides of the metals, e.g,, — 

Zn + 2 HCl = H 2 + ZnCh 

Zinc Hydrochloric Acid Hydrogen Zinc Chloride 

Hydrochloric acid also forms chlorides by interaction with 
(metallic) oxides, hydroxides, and also with carbonates, 
thus:— 

CaO + 2HC1 = CaCb + H 2 O 

Calcium Oxide Calcium Chloride Water 

NaOH + HCl = NaCl + H 2 O 

Sodium Hydroxide Sodium Chloride Water 

CaCOs + 2 HCl = CaCl 2 + CO 2 + H 2 O 

Calcium Hydrochloric Calcium Carbon Water 

Carbonate Acid Chloride Dioxide 

134. Uses of hydrochloric acid. — Hydrochloric acid is 
an indispensable compound, and is used frequently in the 
laboratory and in many industrial processes. Like other 
acids, it is sold in bottles holding 2.5 liters and in securely 
packed glass carboys containing 10 or more gallons (Fig. 60). 

135. Aqua regia. — Hydrochloric acid and nitric acid 
interact and liberate chlorine, thus : — 

3 HCl + HNOs = CI 2 + NOCl + 2 H 2 O 

Hydrochloric Nitric Chlorine Nitrosyl Water 

Acid Acid Chloride 

A mixture of three volumes of concentrated hydrochloric 
acid and one volume of concentrated 
nitric acid is usually used. If such a 
mixture is added to a metal which is 
not affected by either acid alone, the 
free chlorine together with the hydro¬ 
chloric acid converts the metal into a 
soluble compound. Thus, the resist¬ 
ant metal gold forms chlorauric acid 
(HAuCL), which on heating yields sol¬ 
uble gold chloride (AuCL). 

The alchemists named this mixture 
of hydrochloric and nitric acids aqua 
regia, meaning royal water,’’ to emphasize the fact that it 



Fig. 60. — A carboy of 
hydrochloric acid 
packed for shipment. 














110 


A BRIEF COURSE IN CHEMISTRY 


dissolves the “ noble ” metal gold. Another name is nitro- 
hydrochloric acid. 

136. Chlorides. — A chloride is a compound of chlorine 
with another element. Chlorides are formed, as we have 
already seen, by the direct combination of chlorine and 
metals (§ 124) and by the interaction of hydrochloric acid 
with metals, oxides, hydroxides, and carbonates (§ 133). 

Most chlorides are soluble in water. But the chlorides 
of lead (PbCh), silver (AgCl), and one of the chlorides of 
mercury (HgCl) are not soluble. They are formed as in¬ 
soluble solids, when hydrochloric acid, or a soluble chloride, 
is added to a solution of a lead compound, silver compound, 
or the proper mercury compound, e.g., — 

Pb(N 03)2 + 2HC1 = PbCh + 2 HNO 3 

Lead Nitrate Hydrochloric Acid Lead Chloride Nitric Acid 

The formation of insoluble solids by double decomposition 
(and certain other changes) is called precipitation, and the 
solid itself is called a precipitate. Precipitates often have 
properties which can be readily observed or determined. 
Thus, silver chloride is white and curdy, and soon turns 
purple in the light; it dissolves in ammonium hydroxide, 
owing to the formation of a complex soluble compound, 
which, however, is readily transformed by dilute nitric 
acid back into silver chloride. Other chlorides have dif¬ 
ferent properties. Hence, the precipitation and behavior 
of silver chloride serve as a test for hydrochloric acid and 
soluble chlorides. 

137. Names of chlorides. — A molecule of a chloride may 
contain one or more atoms of chlorine and in some cases 
the name of the compound indicates this fact, e.g., man¬ 
ganese dichloride (MnCb), antimony trichloride (SbCh), 
carbon tetrachloride (CCI4). Some metals form two chlo¬ 
rides. Then the two are distinguished by modifying the 
name of the metal; the one containing the smaller propor¬ 
tion of chlorine ends in -ous, that containing the larger in 
-ic. Thus, mercurous chloride is HgCl, and mercuric chlo¬ 
ride is HgCh. 


CHLORINE — HYDROGEN CHLORIDE 


111 


EXERCISES 

1 . How can chlorine be quickly distinguished from other gases? 

2 . Summarize the chemical properties of chlorine. 

3. What is (a) muriatic acid, (b) chloride of lime, (c) bleaching 
powder, (d) carbon tetrachloride, (e) aqua regia ? 

4. Write the equation for (a) preparation of chlorine, (b) interaction 
of chlorine and water, (c) decomposition of hypochlorous acid, (d) prep¬ 
aration of hydrogen chloride, (e) interaction of hydrogen chloride and 
ammonia, (f) interaction of sodium chloride and silver nitrate. 

6. Complete and balance these equations: (a) NaCl -j- H2SO4 = —— 

-f Na 2 S 04 ; (b) H 2 -h- = HCl; (c) HCl -f Na 2 C 03 =-+ CO 2 

+ -; (d) HgNOs +- = HgCl + HNO3. 


PROBLEMS 

1 . Calculate the weight of chlorine in (a) 2 kg. of sodium chloride, 

(b) 2 mg. of calcium chloride, (c) 1 gm. of aluminum chloride. 

2 . How many grams of each product are formed when hydrochloric 
acid interacts with 85 gm. of manganese dioxide? 

3. How many grams of hydrogen chloride can be obtained from 27 
gm. of sodium chloride? 

4. Calculate the percentage composition of (a) KCl, (5) CaCL.' 

6. How much sodium chloride can be formed by burning sodium in 
40 gm. of chlorine? 

6. Calculate the simplest formula of the compound having the com¬ 
position (a) Hg = 84.92 per cent. Cl = 15.07; (b) Hg = 73.8, Cl = 26.2, 

(c) N = 26.17, H = 7.48, Cl = 66.35. 


SUGGESTIONS FOR LABORATORY WORK 


(References are to Newell’s Laboratory Exercises in Chemistry) 


Exercise 

Exercise 

Exercise 

Exercise 

Exercise 

Exercise 

Exercise 

Exercise 

Exercise 

Exercise 


*23 — Preparation and Properties of Chlorine. 

519 — Preparation and Properties of Chlorine — T. 

520 — Chlorine Water — T. 

521 — Bleaching with Bleaching Powder — T. 

*24 — Hydrogen Chlorine and Hydrochloric Acid. 

522 — Hydrogen Chloride. 

*25 — Tests for Hydrochloric Acid and Chlorides. 
*26 — Insoluble Chlorides. 

523 — Aqua Regia — T. 

524 — Types of Chemical Change — T. 





112 


A BRIEF COURSE IN CHEMISTRY 


SUPPLEMENTARY SECTIONS FROM PART II 
The following sections may be selected from Topic XI: — 

363. Gay-Lussac’s law of gas volumes. 

364. Avogadro’s law. 

366. How Avogadro’s law is used to find molecular weights. 

366. Why the molecular weight of oxygen is 32. 

367. Finding molecular weights by the vapor density method. 

368. A molecule of oxygen contains 2 atoms. 

369. Molecules of other elementary gases contain two atoms. 

370. A mole. 

371. How molecular weight is calculated by the molar method. 

372. Molecular formulas of compounds. 

373. Molecular formulas of elements. 

374. Molecular (volumetric) equations. 


CHAPTER XII 


NITROGEN — THE AIR — ARGON AND HELIUM — 
LIQUID AIR 

138. Introduction. — The atmosphere is the gas that 
envelops the earth. It extends several miles into space. 
The terms atmosphere, the air, and air are often used inter¬ 
changeably; though by the air or air we usually mean a 
limited portion of the atmosphere, e.g., the air of a room. 

139. Occurrence of nitrogen. — The gas nitrogen com¬ 
prises about four-fifths by volume (§ 145) of the air. 
There are many compounds of nitrogen, 
e.g., nitric acid (HNO 3 ), sodium nitrate 
(NaNOs), and ammonia (NH 3 ). Nitro¬ 
gen is also a constituent of many animal 
and vegetable substances, e.g., the pro¬ 
teins, which are indispensable ingredients 
of our food and also of the muscles and 
nerves of our bodies. 

140. Preparation of nitrogen. — Nitro¬ 
gen is prepared on an industrial scale 
from liquid air (§§ 155, 156). When 
liquid air is allowed to evaporate slowly, 
the nitrogen escapes first and is collected 
in a separate tank. 

Nitrogen can also be obtained from air by removing the 
oxygen, e.g., industrially from producer gas (§ 65) and in 
the laboratory by phosphorus (Fig. 61). 

Phosphorus is put in a small dish or a crucible cover supported on a 
cork floating in a vessel of water. Upon igniting the phosphorus (Care !) 
with a hot wire and placing a bell jar quickly over the cork, the phos¬ 
phorus and oxygen unite, forming clouds of white phosphorus pentoxidc 

113 



Fig. 61. — Prepara¬ 
tion of nitrogen 
from air by burn¬ 
ing out the oxygen 
with phosphorus. 









114 


A BRIEF COURSE IN CHEMISTRY 


(P2O5). This sohd soon dissolves in the water, which rises inside the 
jar owing to the removal of the oxygen, and the nitrogen is finally left. 
Caution : This experiment is dangerous and great care must be used. 

The most convenient way to prepare nitrogen in the 
laboratory is by heating a solution of sodium nitrite (NaN 02 ) 
and ammonium chloride (NH4CI). These two compounds 
form the unstable compound ammonium nitrite (NH4NO2), 
which decomposes into nitrogen and water, thus: — 

NH4NO2 = N2 + H2O 

Ammonium nitrite Nitrogen Water 

Small quantities of nitrogen may be readily obtained 
by heating ammonium dichromate ((NH 4 ) 2 Cr 207 ). 

141. Physical and chemical properties of nitrogen. — 
Nitrogen is a colorless, odorless gas. It is a little lighter 
than oxygen and air. A liter weighs 1.25 gm. (A liter of 
oxygen weighs 1.43 gm. and one of air 1.29 gm.) It is only 
slightly soluble in water. 

Nitrogen does not support combustion nor sustain life. 
Flames are extinguished by nitrogen and animals are suffo¬ 
cated by it. 

Nitrogen is very much less active chemically than oxygen. 
Indeed, it responds to none of the common tests. It is some¬ 
times called an inert element. At high temperatures and 
under special conditions, however, nitrogen is active and 
forms many compounds. Thus, it combines with magnesium 
and a few other metals at red heat, forming nitrides, e.g., 
magnesium nitride (Mg 3 N 2 ). An electric current causes 
nitrogen to combine with oxygen and with hydrogen, forming 
nitrogen oxides (NO, NO2) and ammonia (NH3). Both 
reactions, if hastened by a catalyst, proceed rapidly and are 
utilized on an industrial scale to convert nitrogen from the 
air into compounds needed as fertilizers and explosives. 
The process of converting nitrogen gas into compounds is 
called “ fixation of nitrogen.’’ 

142. Uses of nitrogen. — Nitrogen on account of its 
inertness is used to fill some kinds of electric light bulbs, 
and the stem of high-boiling thermometers. It is also used 


NITROGEN —THE AIR —ARGON AND HELIUM 115 


in making ammonia, nitric acid, and a nitrogen fertilizer 
called calcium cyanamide (CaCN2) (§ 143). 

143. Nitrogen and life. — Nitrogen, as well as oxygen, 
is vitally connected with life, though in a different way. 
All animals need nitrogen for their growth. But although 
we live in an atmosphere containing nearly 80 per cent 
of this gas, we can not assimilate it directly. The nitrogen 
we inhale (along with the oxygen) is exhaled again unused. 
The nitrogen needed by our bodies must be eaten in the form 
of nitrogenous food, such as lean meat, fish, and wheat. 

Nor have plants, with few exceptions (see next paragraph) 
power to assimilate free nitrogen from the atmosphere. 
Most plants take up com¬ 
bined nitrogen from the 
soil in the form of nitrates 
or of ammonia. Hence 
combined nitrogen is be¬ 
ing constantly removed 
from the soil. In order to 
restore it, a nitrogen com¬ 
pound must be added, 
e.g., sodium nitrate 
(NaNOs), calcium nitrate 
(Ca(N03)2), ammonium 
chloride (NH4CI), ammo¬ 
nium sulfate ((NH4)2S04), 
or calcium cyanamide 
(CaCN2). A substance 
or mixture which restores nitrogen (or some other chemical 
element like phosphorus or potassium) to the soil is called a 
fertilizer. 

Leguminous plants, such as peas, beans, and clover, get 
nitrogen from the air (in the loose soil) by means of bacteria, 
which are in nodules on their roots (Fig. 62 ). The bacteria 
change the nitrogen into compounds which can be utilized 
by the growing plant. Sometimes this process is also called 

fixation of nitrogen.’’ 

144. Air is a mixture and not a compound. Air is a 

mixture of several gases. Oxygen, nitrogen, and argon 








116 


A BRIEF COURSE IN CHEMISTRY 


are the three ingredients that are always present in nearly 
constant proportions. Variable proportions of water vapor 
and carbon dioxide are always found, and also small quan¬ 
tities of compounds related to ammonia and nitric acid. 
Near cities the air may contain considerable dust, sulfur 
compounds, and acids; at the ocean some salt is found. 

The following facts show that air is not a compound but 
a mixture of gases: — 

(1) The proportion of oxygen and of nitrogen is not fixed 
but varies between small limits. Therefore air does not 
have a constant composition and can not be represented 
by a formula. 

(2) When nitrogen and oxygen are mixed in approximately 
the proportions that form air, the product is identical with 
air, but the act of mixing gives no evidence of chemical 
action, e.g., no heat or light is produced. 

(3) When air is dissolved in water, a larger proportion of 
oxygen than nitrogen dissolves. If the oxygen and nitrogen 
were combined, the dissolved air would contain the same 
proportions of oxygen and nitrogen as the original air. 

(4) When air is liquefied and allowed to evaporate, the 
nitrogen escapes first (§ 156). If air were a compound, 
liquid air would evaporate as a whole. 

145. Proportions of the main ingredients of air. — The 
normal proportions (by volume) are nitrogen 78.06 per cent, 
oxygen 21, and argon 0.937. These numbers are often 
rounded off to nitrogen 78, oxygen 21, and argon 0.94. 

146. ** Composition ” of air. — Air near the surface of 
various parts of the globe shows such uniformity in the 
proportions of the main ingredients that chemists have 
fallen into the habit of applying the term composition to air. 

The proportion of oxygen in the air can be readily found in the labo¬ 
ratory by two methods. (1) A known volume of air is shaken in a 
closed bottle with a mixture of pyrogallic acid and sodium hydroxide; 
this solution absorbs the oxygen and leaves the nitrogen and argon un¬ 
changed. (2) A graduated glass tube, containing a known volume of 
air, is inverted in a jar of water, and a piece of phosphorus attached to a 
wire is pushed up into the tube (Fig. 63). The oxygen combines with 
the phosphorus. In a few hours the phosphorus is removed, and the 
volume of residual gas is read. 


NITROGEN —THE AIR —ARGON AND HELIUM 117 

In each process the difference between the first and last volumes is 
oxygen. There is no simple way of separating the nitrogen and argon. 

147. Water vapor in the air. — Water vapor is always 
present in air, owing to constant evaporation from the ocean 
and other bodies of water. The per cent in a given place 
depends on the temperature, though it is influenced also 
by pressure, winds, and configuration of the land. When 
the temperature of the air falls sufficiently, the water vapor 
condenses and is deposited in the form of dew, rain, fog, mist, 
frost, snow, sleet, or hail. The clouds are masses of minute 
drops of liquid water formed by condensation of the water 
vapor in the cold upper air. The conden 
sation of considerable moisture forms large 
drops, which fall as rain. 

The proportion of water vapor in the air 
of different regions varies between wide 
limits. Thus, in tropical countries it is 
large, while in desert countries it is small. 

148. Test for water vapor in the air. — 

Moisture collects on the outside of a vessel 
containing cold water, such as a pitcher of 
iced water, on water pipes in a cellar. 

The presence of water vapor may be 
shown also by exposing to the air a deli¬ 
quescent substance, such as calcium chlo¬ 
ride, which soon becomes moist (§ 94). 

149. Physical comfort depends on water 
vapor in the air. — Our bodies have a 
normal and nearly constant temperature 
of 37° C. (98.6° F.). This temperature is 
maintained by the heat produced by the 
chemical changes in our bodies, especially 
the oxidation of waste tissue by the oxygen carried by the 
blood to all parts of the body (§34). This temperature is 
regulated partly by radiation of heat and partly by evap¬ 
oration of water from the surface of the body. In moist air, 
evaporation proceeds too slowly; in dry air, too fast. In 
either case, we are uncomfortable, and try in various ways 
to become comfortable. Thus, we use fans and wear thin 



of air by phos¬ 
phorus. 


























118 


A BRIEF COURSE IN CHEMISTRY 


clothing to promote evaporation, or we moisten the air by 
exposing pans of water. 

160. Carbon dioxide in the air. — Carbon dioxide is one 
product of the breathing of animals, the combustion of fuels, 
and the decay of vegetable and animal substances (§ 64). 
By these processes vast quantities of carbon dioxide are 
being constantly introduced into the air. Thus, an average 
person exhales about 1000 gm. of carbon dioxide daily. 

The proportion of carbon dioxide in ordinary air is 3 to 4 
parts in 10,000 parts of air, ^.e., 0.03 to 0.04 per cent. In 
crowded rooms it may be as high as 1 per cent or even higher 
in unusual cases, e.gr., a crowded, badly ventilated room. 

The proportion of carbon dioxide in the atmosphere as a 
whole is practically constant, owing to the mixing by winds 
and air currents, and largely, also, to absorption of this gas 
by all green plants •(§ 69). 

161. Test for carbon dioxide in the air. — Carbon dioxide 
in the air interacts with calcium hydroxide, forming a thin, 
white crust of insoluble calcium carbonate on the surface of 
the liquid. If considerable air is drawn through calcium 
hydroxide solution, the liquid becomes milky, because the 
particles of calcium carbonate are suspended in the liquid 
(§64 and Fig. 19). 

162. Argon in the air. — Argon is an essential and constant 
ingredient of the air, the proportion being 0.937 per cent. 

Argon is a colorless, odorless gas which is a little heavier 
than oxygen. It dissolves in water to the extent of about 
4 volumes in 100. 

A conspicuous property of argon is its utter lack of chemical 
activity. No compounds of this element have yet been 
prepared or discovered. On account of its inertness, argon 
is used to fill one kind of electric light bulb. 

163. Other gases in the atmosphere. — Helium, neon, 
krypton, and xenon are inert gases discovered by Ramsay. 
With the exception of neon, they constitute an exceedingly 
minute proportion of the atmosphere. Like argon they do 
not form compounds. Helium is used to inflate dirigible 
airships and neon is used to produce the red illuminated 
advertising signs. 


NITROGEN —THE AIR —ARGON AND HELIUM 119 


154. Helium. — Helium is obtained in large quantities 
from natural gas which issues from the ground in certain 
parts of the United States, e.g., Kansas, Oklahoma, and 
Texas. The gas contains about 2.5 per cent of helium which 
is extracted on a large scale by a complex process and com¬ 
pressed into large tanks for use in airships (Fig. 64). It is 
a light, non-combustible gas. Both properties make helium 
an excellent gas for filling balloons and dirigible airships. 
Helium is now used instead of hydrogen as the lifting gas in 
the United States airships (Fig. 65). 

155. Liquid air. — Liquid air is essentially a mixture of 
liquid oxygen and liquid nitrogen. It sometimes looks 



Fig. 64. — Car containing three tanks of compressed helium. 


eloudy, owing to the presence of solid carbon dioxide and 
ice. When these solids are removed by filtering, the liquid 
air has a pale blue tint. 

156. Properties of liquid air. — If a beaker is filled with 
liquid air, the latter boils vigorously, the surrounding gaseous 
air becomes intensely cold, frost gathers on the beaker, and 
in a short time the liquid air will disappear into the air of the 
room. If, however, liquid air is put into a Dewar flask, 
evaporation takes place so slowly that some liquid air will 
remain in the flask several days. 

A Dewar flask consists of two flasks, one within the other sealed 
together air-tight at the top; the space between the flasks is a vacuum. 





120 


A BRIEF COURSE IN CHEMISTRY 


The surfaces of the flasks are coated with silver, which reflects heat and 
helps retard the evaporation of the hquid air. Liquid air is stored and 
transported in Dewar flasks. Thermos bottles are constructed like 
Dewar flasks. 

Liquid air is a mixture and does not have a fixed boiling 
point, though fresh samples boil at about —190° C. If it 




Fig. 65. — The Los Angeles buoyed in the air by bags of helium. 


is allowed to boil in a proper apparatus, the nitrogen (boiling 
point — 194°C.) escapes first, leaving more or less pure 
oxygen (boiling point —182.5° C.). The industrial sepa¬ 
ration is accomplished this way, 
the two gases being collected in 
separate tanks. 

167. Some experiments with liquid 
air. — A tin or iron vessel which has 
been cooled by liquid air is so brittle 
that it may often be crushed with the 
fingers; while a piece of rubber tubing 
becomes as brittle as glass. Mercury 
freezes so hard in liquid air that it can 
be used as a hammer to drive a nail. 

When liquid air is poured into a ves¬ 
sel standing on a cake of ice (Fig. 66), 
the liquid air boils vigorously because 
the ice is so much “ hotter.” If a kettle 
of liquid air is placed over a lighted Bunsen burner, frost and ice collect 
on the bottom of the kettle, because the intense cold produced by the 



Fig. 66. — Vessel of liquid air 
boiling on a cake of ice. 









il 


NITROGEN —THE AIR —ARGON AND HELIUM 121 

evaporation of the liquid air in the kettle solidifies the water vapor 
and carbon dioxide, which are the two main products of burning illu¬ 
minating gas. If water is now poured into the kettle, the liquid air 
boils over and the water is instantly frozen; the water is so much 
“ hotter ” than the liquid air that the latter boils more violently, and 
since its rapid evaporation causes absorption of heat, the water loses 
heat and becomes ice. 

Ordinary liquid air is from one-half to one-fifth liquid oxygen, and 
will support combustion. A glowing stick or a red-hot rod of steel burns 
brilliantly in this cold liquid. 

158. Manufacture of liquid air. — Liquid air is manu¬ 
factured by subjecting air to a high pressure and cooling it 



Fig. 67. — Diagram of apparatus for liquefying air. 


to a low temperature. A diagram of the essential parts of 
a liquid air machine is shown in Fig. 67. 

Air freed from moisture and carbon dioxide is forced through the 
valves A and B into the compressor C. Here a pump compresses the 
air and forces it out through the valve D into the coiled pipe FF where 
the heat (produced by the compression) is absorbed by running water. 
The compressed and cooled air is then forced along through the inner 
of two pipes to the valve H. This valve is small and the air escapes 
through it into a chamber where the pressure is lower (e. g., 1 atmosphere). 
As the air passes through the fine opening, it expands rapidly and be¬ 
comes very cold. This cold gas is made to flow back through the outer 
pipe G, and in so doing cools the gas flowing forward in the inner pipe. 

















































122 


A BRIEF COURSE IN CHEMISTRY 


The backward flowing gas passes through the pipe I, enters the machine 
again at B, is compressed, cooled, and expanded — as before. This 
process is continued until finally the air in the inner pipe is cooled so 
low that it liquefies in part, and drops from H into the liquid air con¬ 
tainer, from which it can be drawn off as needed through the valve at J. 

EXERCISES 

1. How is nitrogen prepared? Summarize its physical properties. 
Compare the chemical properties of nitrogen and oxygen. 

2. What is the relation of nitrogen to the life of (a) animals and (b) 
plants ? 

3. What are the two chief ingredients of the atmosphere ? The con¬ 
stant ingredients? The variable ingredients? The ingredients found 
in traces? 

4. State the volumetric composition of air. How is it found? 

6. Give three proofs that air is a mixture. 

6. (a) How would you distinguish nitrogen from carbon dioxide? 
(b) What two properties of helium make it useful in airships? 

PROBLEMS 

(See also Problems at the end of Chapters VIII and IX.) 

1. What is the weight of air in a room 6X8X5m.? (A liter of air 
weighs 1.29 gm.) 

2. How many kilograms of pure, dry air are needed to yield (a) 100 
kg. and (b) 100 1. of oxygen (at standard conditions) ? 

SUGGESTIONS FOR LABORATORY WORK 
(References are to Newell’s Laboratory Exercises in Chemistry) 

Exercise 27 — Per cent of Oxygen in Air — T. 

Exercise S25 — Air and Combustion. 

Exercise S26 — Water Vapor and Carbon Dioxide in Air. 

Exercise S27 — Preparation and Properties of Nitrogen — T. 

Exercise 4 — Heating a Known Weight of Metal in Air. 

SUPPLEMENTARY SECTIONS FROM PART II 

376. Atomic weights. 

376. Finding approximate atomic weights. 

377. Exact atomic weights. 

378. International atomic weights. 


CHAPTER XIII 


ACIDS, SALTS, AND BASES 

169. Acids. — Solutions of acids (1) have a sour taste, 
(2) turn blue litmus red, (3) interact with certain metals, 

(4) interact with oxides, hydroxides, and carbonates, and 

(5) conduct electricity. 

Hydrogen is an essential constituent of all acids. The 
other constituent is a non-metal or a group of non-metals, 
e.g., Cl, SO4. The hydrogen in acids can be liberated by 
certain metals, and the compound formed by this replace¬ 
ment is called a salt. Many compounds have hydrogen as 
a constituent, but they are not acids unless they form salts 
by replacement of the hydrogen by a metal. 

A simple test for acids is the change they produce in the 
color of litmus — blue to red. Acids cause other substances 
to change color, e.g., phenol-phthalein from pink to colorless. 
Such substances are used to detect acids and are called 
indicators. Acids and other substances which redden blue 
litmus, or produce a specific color change in an indicator, 
are said to have an acid reaction. The presence of acids is 
often conveniently shown by the litmus test. 

Common acids are hydrochloric acid (HCl), sulfuric 
(H2SO4), nitric (HNO3), and acetic (HC2H3O2). 

160. Salts. — This class of compounds has many members 
and their properties vary greatly. (1) Many salts have the 
“ salty ” taste associated with common salt (sodium chloride). 
However, some are sour, others are bitter, and a few are 
tasteless. (2) Their solutions do not behave alike with 
litmus. Some have no effect on litmus and other indicators, 
and are said to be neutral or to have a neutral reaction. But 

123 


124 


A BRIEF COURSE IN CHEMISTRY 


some turn blue litmus red, while others act oppositely. 
(3) Solutions of salts conduct electricity (§ 164). 

Salts invariably have a metal and a non-metal as constitu¬ 
ents, and most salts also have oxygen. Chlorides are ex¬ 
amples of salts which have only a metal and the non-metal 
chlorine as constituents, e.g., sodium chloride (NaCl), 
calcium chloride (CaCh). Examples of salts composed of a 
metal, a non-metal, and also oxygen are potassium chlorate 
(KCIO3) and sodium sulfate (Na 2 S 04 ); these two com¬ 
pounds are salts of chloric acid and sulfuric acid respectively. 

161. Bases. — Hydroxides are examples of the class of 
compounds called bases. Thus, sodium hydroxide (NaOH) 
is a typical base. The general properties of bases differ 
from those of acids and salts. (1) Solutions of bases turn 
red litmus blue — the opposite of acids — and are said to have 
a basic or an alkaline reaction. Bases also turn phenol- 
phthalein solution from colorless to rose-pink — the opposite 
of acids. These color changes serve as a test for bases. 
(2) Solutions of strong bases {e.g., sodium hydroxide and 
potassium hydroxide) have a slippery feeling and a biting, 
caustic taste. Solutions of bases (3) interact with acids 
and form salts and water, and (4) conduct electricity. 

A base is composed of a metal, combined with the group 
OH called hydroxyl, e.g.y sodium hydroxide is NaOH. 
This group of oxygen and hydrogen is called a radical, and is 
the essential part of a base, just as the hydrogen is of an acid. 
In fact, the oxygen and hydrogen of a base act as a unit in 
many chemical changes. 

162. Neutralization. — When we mix a solution of an acid 
with a solution of a base in the proper proportions, the acid 
and base interact completely. The final solution has none 
of the characteristic properties of either an acid or a base, 
but it does have the properties of a salt. That is, the acid 
and base neutralize each other. For example, hydrochloric 
acid and sodium hydroxide interact, and form sodium 
chloride and water, thus: — 


Acid 


Base 


Salt Water 


HCl + NaOH 


NaCl + H 2 O 


Hydrochloric Sodium 


Sodium 

Hydroxide 


Sodium Water 

Chloride 


Acid 


ACIDS, SALTS, AND BASES 


125 


A chemical change in which an acid and a base neutralize 
each other and thereby form a salt and water is called 

neutralization. 

Neutralization illustrates double decomposition, i.e., a 
decomposition into parts and a recombination on another 
plan. In the chemical change just cited both the hydro¬ 
chloric acid and the sodium hydroxide are decomposed and 
their parts are recombined into sodium chloride and water. 

Neutralization may be accomplished by carefully mixing an acid and 
a base. But it is usually done by burettes. These are graduated glass 
tubes, so marked that any desired portion of the 
contents can be drawn off at the lower end (Fig. 

68). In using burettes for neutralization, one is 
fiUed to the zero (top) mark with a solution of an 
acid and the other with a solution of a base — one 
(sometimes each) solution being of known con¬ 
centration. A measured portion, say 15 cubic 
centimeters of the base, is drawn off into a beaker, 
several drops of litmus (or phenol-phthalein) solu¬ 
tion are added, and the acid is slowly dropped in 
from the burette with constant stirring until one 
drop more shows by the change in color (after 
thorough stirring) that the right proportions of « 
acid and base are present, i.e., that neutralization 
has occurred. c 

If we wish to find the strength of the acid, the 
volume of acid is read accurately. Knowing the y ^ 
concentration (and volume) of the base solution ^ 
and the volume of the acid solution, we can calcu- Fig. 68. — Burettes, 
late the exact weight of the acid needed for the 

neutrahzation of the base, and from this weight we can find the strength 
of the acid solution. Suppose in adding hydrochloric acid to a definite 
volume of sodium hydroxide solution of known strength, it was found 
that 1 cc. of the sodium hydroxide solution equals 1.5 cc. of the hydro¬ 
chloric acid solution. The problem is to calculate the number of grams 
of HCl in 1 cc. of the solution of hydrochloric acid used in this experi¬ 
ment. Proceed as follows: 

(а) First write the equation for the reaction, thus: — 

HCl -h NaOH = NaCl + H 2 O 
36.5 40 58.5 18 

This equation means that 36.5 gm. of HCl are needed to neutrahze 40 
gm. of NaOH. 

(б) Next find the average number of cubic centimeters of hydro¬ 
chloric acid solution neutralized by 1 cc. of sodium hydroxide solution. 






















126 


A BRIEF COURSE IN CHEMISTRY 


In this illustration we assume that 1.5 cc. HCl sol. = 1 cc. NaOH sol. 
(Your result, of course, may be different from this value.) 

(c) Next learn from the Teacher the concentration of the sodium 
hydroxide solution. Suppose 1 cc. of sodium hydroxide solution con¬ 
tains 0.00641 gm. of NaOH. 

(d) Now from the equation in (a) 40 gm. of NaOH require 36.5 gm. 
of HCl. Then the number of grams of HCl required by 0.00641 gm. of 
NaOH would be found by the proportion 

40: 36.5:: 0.00641: x a: = 0.00585 

(Your result depends on the concentration of your NaOH solution.) 
But 0.00585 gm. of HCl would be dissolved in 1.5 cc. of hydrochloric 
acid (according to our supposition in (6)). Therefore, to find the number 
of grams of HCl that would be dissolved in 1 cc. of the acid solution, 
we divide 0.00585 by 1.5, ^.e., 0.00585 ^ 1.5 = 0.0039. Ans. 0.0039 
gm. of HCl in 1 cc. 

163. Another definition of a salt. — For the present, we 
may regard a salt as a compound formed from an acid and a 

base by neutralization. 
That is, the metal of 
the base unites with 
the non-metal or non- 
metallic group of the 
acid, e.g., Na of NaOH 
unites with NO 3 of 
HNOsto formNaNOs. 

164. Solutions of 
acids, salts, and bases 
conduct electricity. — 
This is a striking and 
fundamental property 
of acids, salts, and bases. It can be demonstrated by the 
apparatus shown in Fig. 69. 

The beaker, or jar, A contains the solution into which dip two pieces 
of platinum, B and C, called the electrodes; through them the current 
enters and leaves the solution. The wire from the electrode C and the 
wire E from the bulb D are connected with the (direct) fight current, 
or storage battery. The wire from the electrode B is also connected 
with the electric fight bulb D. When the (direct) current is turned on, 
the bulb glows if the solution contains an acid, salt, or base. For 
example, if solutions of hydrochloric acid, calcium chloride, and sodium 
hydroxide are tried in succession, the bulb glows brightly in each case. 



Fig. 69. — Apparatus for showing that only 
solutions of acids, bases, and salts conduct 
an electric current. 















ACIDS, SALTS, AND BASES 


127 


165. Naming acids. — The scientific names of acids are 

based on their composition. 

The few acids containing only two elements have the 
prefix hydro- and the suffix -ic coupled with a simple modifica¬ 
tion of the name of the non-metallic element. Thus, bromine 
forms hydrobromic acid (HBr); chlorine, hydrochloric 
(HCl); fluorine, hydrofluoric (HF); iodine, hydriodic (HI) ; 
and sulfur, hydrosulfuric (H 2 S). Remember these names. 

Most acids contain oxygen, and their names are based on 
its proportion. The common acid of an element has the 
suffix -ic, and the acid containing one less atom of oxygen 
has the suffix -ous. Thus, nitric acid is HNO3 and nitrous 
acid is HNO2, sulfuric acid is H2SO4 and sulfurous acid is 
H2SO3. If an element forms an acid containing less oxygen 
than the -ous acid, the name of this acid has in addition the 
prefix hypo- (meaning less); e.g., hypochlorous acid (HCIO). 
If an element forms an acid containing more oxygen than the 
-ic acid, such an acid retains the suffix -ic, and has in addition 
the prefix per- (meaning more), e.g., perchloric acid (HCIO4). 
A confusing point in learning the names of acids is the modi¬ 
fication of the name of the ' characteristic element. These 
must be learned from the names of acids met in studying, 
e.g., the element phosphorus forms phosphoric acid, and 
silicon forms silicic acid. 

166. Naming salts. — Salts are named from their corre¬ 
sponding acids. In the names of salts containing only two 
elements or groups, the sufllx -ide replaces -ic of the acid, 
giving the names chloride, bromide, sulfide, fluoride, and 
iodide. The prefix hydro- in the name of a binary acid is 
omitted from the name of the salt. Thus, the sodium salt 
of hydrochloric acid is sodium chloride, not hydrochloride; 
similarly, there are potassium bromide, lead sulfide, calcium 
fluoride, and sodium iodide. Two salts containing the same 
elements are distinguished by adding -ous and -ic to a slight 
modification of the name (English or Latin) of the metal — 
-ous to the one containing the smaller per cent of the non- 
metal. Thus, HgCl is mercurous chloride and HgCh is 
mercuric chloride. 

Most salts contain oxygen. Their names are derived 


128 


A BRIEF COURSE IN CHEMISTRY 


from the names of the corresponding acids by changing the 
suffixes -ic to -aie and -ou8 to -ite; the particular salt is of 
course distinguished by the name of the metal. The prefixes 
and per- are retained. The following are examples: — 


Acid 

Hypochlorous acid . 
Sulfurous acid. . . 

Phosphoric acid . . 

Permanganic acid 


Corresponding Salt 
HCIO Sodium hypochlorite . NaClO 

H2SO3 Potassium sulfite . . K2SO3 

H3PO4 Calcium phosphate . . Ca3(PO)2 

HMn04 Silver permanganate . AgMn04 


Note that the modification of the name of some elements 
differs in corresponding acids and salts — sulfur- and sulf-, 
phosphor- and phosph-. 

167. Naming bases. — Bases are named by placing the 
name of the metal (or group) before the word hydroxide, 
6.gf., sodium hydroxide (NaOH). 

168. Formation of acids and bases from oxides and water. 

— We learned in § 85 that water combines directly with 
many oxides, e.gr., sulfur dioxide and calcium oxide, thus 

SO2 + H2O = H2SO3 

Sulfur Dioxide Water ' Sulfurous Acid 

CaO + H 2 O = Ca(OH )2 

Calcium Oxide Water Calcium Hydroxide 

Many oxides act similarly. The oxides of non-metals, such 
as carbon, sulfur, and phosphorus, produce acids. Whereas 
the oxides of metals, such as sodium, calcium, and mag¬ 
nesium, produce bases. Such oxides are called anhydrides 

— acid (or acidic) anhydrides if they form acids, and basic 
if they form bases. 

169. Formation of salts. — Several methods have already 
been given. Let us review them and add others. 

A. The following methods give salts soluble in water: — 

(a) An acid and a metal — the usual method (§ 68), e.gr., 
zinc and sulfuric acid form zinc sulfate, and calcium and 
hydrochloric acid form calcium chloride. We can show a 
salt is formed by evaporating the solution to dryness and 
testing small portions of the residue for the parts of the salt, 
Le., in these two cases, for the metal part (zinc and calcium) 
and the non-metal part (sulfate and chloride). 


ACIDS, SALTS, AND BASES 129 

(6) An acid and an oxide, e.gr., hydrochloric acid and 
calcium oxide. 

(c) An acid and a carbonate, e.g.j hydrochloric acid and 
calcium carbonate. 

(d) An acid and a base — neutralization, e.g., hydrochloric 
acid and sodium hydroxide. 

B. The following methods give salts insoluble in water : — 

(a) An acid and a salt, e.g., sulfuric acid and lead nitrate 
form insoluble lead sulfate (and nitric acid). Similarly 
sulfuric acid and barium chloride form insoluble barium sul¬ 
fate (and hydrochloric acid). 

(h) A salt and a salt, e.g., sodium sulfate and lead nitrate 
form insoluble lead sulfate (and soluble sodium nitrate). 
Similarly, sodium chloride and silver nitrate form insoluble 
silver chloride (and soluble sodium nitrate). 

(c) An acid and an oxide, e.g., dilute hydrochloric acid 
and lead oxide (PbO) form insoluble (in cold water) lead 
chloride. 

(d) An acid and a carbonate, e.g., hydrochloric acid and 
lead carbonate form insoluble (in cold water) lead chloride. 

One kind of salt called a normal salt is formed by neutral¬ 
izing exactly the proper pair of acid and base. Thus, sulfuric 
acid and sodium hydroxide form normal sodium sulfate 
(Na 2 S 04 ) by exact neutralization. But if the acid and base 
are neutralized and then more sulfuric acid is added, the 
crystals obtained by evaporating the solution are those of 
acid sodium sulfate (NaHS 04 ). Analogous salts are normal 
sodium carbonate (Na 2 C 03 ) and acid sodium carbonate 
(NaHCOs), and normal calcium carbonate (CaCOs) and 
acid calcium carbonate (CaH 2 (C 03 ) 2 ). 

EXERCISES 

1. State two characteristics of (a) acids, (&) bases, (c) salts. 

2. Give the name and formula of three common (a) acids, (b) bases, 
and (c) salts. 

3. Define and illustrate neutralization. Write an equation. 

4. Write the equations for the reactions in A and B (§ 169). 

6. Give the name and formula of the sodium salt of hydrochloric 
acid. Also of the salt corresponding to potassium, aluminum, lead, 
silver, manganese, zinc, and barium. 


130 


A BRIEF COURSE IN CHEMISTRY 


6. Apply Exercise 6 to (a) nitric acid, (&) nitrous acid, (c) hypo- 
chlorous acid. 

7. Apply Exercise 6 to (o) sulfuric acid, ( 6 ) sulfurous acid, (c) per¬ 
manganic acid. 

8 . Give the name and formula of the hydroxide of the metals enu¬ 
merated in Exercise 5. 

9. Give the name of these : (a) potassium salt of chloric acid, (b) cal¬ 
cium salt of hypophosphorous acid, (c) sodium salt of carbonic acid, 
(d) lead salt of chromic acid, (e) zinc salt of hydriodic acid, (f) potassium 
salt of perchloric acid, (gr) iron salt of hydrochloric acid, (h) sodium salt 
of hydrofluoric acid, (i) calcium salt of persulfuric acid, (j) potassium 
salt of hydrobromic acid, (k) calcium salt of hydrofluoric acid, (1) sodium 
salt of hypophosphorous acid. 

10. Complete and balance: (a) BaO-|-= Ba(OH) 2 ; ( 6 ) NHJ 

+- = Agl -f (NH4)2S04; (c) Pb(N03)2 +- = PbCl 2 +-. 

PROBLEMS 

1. Calculate the per cent of hydrogen in (a) hydrochloric acid, 
(b) sulfuric acid, (c) nitric acid. 

2. Calculate the per cent of hydroxyl in (a) sodium hydroxide, 
( 6 ) potassium hydroxide, (c) ammonium hydroxide. 

3. What weight of HNO 3 will neutralize 27 gm. of NaOH ? 

4. Calculate the formula corresponding to: (a) Ca = 29.41, 

S = 23.52, O = 47.05; (b) Na = 39.31, Cl = 60.68. 

6. What weight of the salt is formed in these cases of neutralization ? 
(a) Hydrochloric acid and 10 gm. of potassium hydroxide; (b) sulfuric 
acid and 37 gm. of sodium hydroxide. 

SUGGESTIONS FOR LABORATORY WORK 

(References are to Newell’s Laboratory Exercises in Chemistry) 

Exercise S28 — General Properties of Acids. 

Exercise S29 — General Properties of Bases. 

Exercise S30 — Two Properties of Many Salts. 

Exercise *28 — Behavior of Oxides with Water. 

Exercise 29 — Neutralization. 

Exercise 30 — Neutralization by Titration — T. 

Exercise 31 — Preparation of Salts. 




CHAPTER XIV 

IONS AND IONIZATION 

170. The ionic theory. — The properties of dilute solu¬ 
tions of acids, bases, and salts are best explained by the 
ionic theory. This theory 
was proposed in 1887 by 
the Swedish chemist 
Arrhenius (Fig. 70). 

The theory is usually 
stated as follows : — 

Acids, bases, and salts, 
when dissolved in water, 
break up into their parts, 
each part being charged 
with electricity. 

This theory means that 
a solution of sodium chlo¬ 
ride, for example, contains 
particles of electrically 
charged sodium and elec¬ 
trically charged chlorine. 

171. What are ions? 

— The breaking up, or Fig. 70. — The Swedish chemist Arrhenius 

dissociation as it is often (1869-1927), who gave the erst explana- 
,, , - . - . tion of the properties oi acids, bases, 

called, of acids, bases, and salts, 

and salts in solution is 

called ionization. The electrically charged parts are called 
ions. Each ion is a portion of a molecule. Two kinds of 
ions are present in every electrolytic solution, viz., electro¬ 
positive ions, or cations, and electro-negative ions, or anions. 

• 131 





132 


A BRIEF COURSE IN CHEMISTRY 


Ions, although formed by the dissociation of molecules, 
are not identical with atoms. Ions are electrically charged 
atoms or radicals. For example, in a solution of sodium 
chloride the electro-positive sodium ions move about in the 
water without producing any hydrogen; whereas ordinary 
sodium interacts violently with water and produces hydrogen, 
as we have already seen (§ 69). Similarly, the chlorine 
ions circulate freely in water without escaping as chlorine 
gas or producing chlorine water. It should be distinctly 
understood that the electric charges on the ions in a solu¬ 
tion do not come from the electricity that may subsequently 
be passed through the solution. The parts of the molecules 
become electrically charged just as soon as the molecules 
break up in the solution. 

172. How ions are represented. — Ions are represented 
by chemical symbols supplemented by a plus (+) or a minus 
(—) sign, which shows the kind of electric charge. Thus, 
the ions formed by sodium chloride are Na+ and Cl“ and 
from copper sulfate are Cu++ and SO 4 . On the other 

hand in calcium chloride (CaCh) solution, each molecule 
dissociates into one calcium ion and two chlorine ions, 
^.e., Ca++ and 2 Cl“ {not Cb"). 

The solution as a whole is electrically neutral. This 
fact means that in the solution the sum of each kind of 
electric charge must be the same (^.e., the total positive 
charge equals the total negative). It does not mean, how¬ 
ever, that the solution must contain the same number of 
positive as negative ions. Thus, the ions formed by calcium 
hydroxide (Ca(OH) 2 ) are Ca++ and 2 0H“; the positive 
charge is 2 and the total negative is 2 (^.c., 2 X 1). In the 
formulas of ions, the coefficients multiply the charges. Thus, 
aluminum sulfate (Al2(S04)3) forms 2 A 1 +++ and 3 SO4—, 
making + 6 and — 6 . 

In ordinary chemical formulas, atoms and radicals are 
represented as united, e.g. CaCb, Ca(OH) 2 . But when the 
molecule dissociates into independent particles (ions), we 
represent the ions as separate particles. That is, the CI 2 
in CaCb becomes 2 Cl“ or Cl~ + Cl~, not Cb”; similarly 
from Ca(OH )2 we have 2 OH~ {not OH 2 “). 


IONS AND IONIZATION 


133 


173. Acid and base defined in terms of ions. — Specific 
properties are exhibited by solutions of typical acids and 
bases. Thus, acids have a sour taste and turn litmus red; 
bases have a bitter taste and turn litmus blue. Solutions 
of acids contain hydrogen ions, and of bases contain hydroxyl 
ions. The properties of such solutions are doubtless largely 
due to the ions. According to the ionic theory, then, an 
acid is a compound whose solution contains hydrogen ions 
(H+), while a base is a compound whose solution contains 
hydroxyl ions (OH~). These definitions should be compared 
with those previously given (§§ 159, 161). 

174. Salts and ionization. — Salts have already been 
defined in several ways. They are compounds (other than 
water) resulting from (1) the neutralization of acids and 
bases, (2) the substitution of a metal for the hydrogen of an 
acid, and (3) the substitution of a non-metal for the hydroxyl 
of a base (§ 163). 

In terms of the ionic theory, salts are compounds which 
in solution yield neither hydrogen nor hydroxyl ions, but 
instead the positive ion of a base and the negative ion of an 
acid. Thus, sodium nitrate yields Na+, which is the posi¬ 
tive ion from the base NaOH, and NOa', which is the negative 
ion of nitric acid (HNO 3 ). An acid salt (§ 169) yields posi¬ 
tive and negative ions, e.gr., NaHS 04 yields Na+ and HS 04 “. 

175. Neutralization defined in terms of ions. — In § 162 
we learned that neutralization is a chemical change in which 
an acid and a base form a salt and water. Neutralization, 
interpreted by the ionic theory, is the combining of hydrogen 
and hydroxyl ions to form molecules of water. Suppose 
solutions of hydrochloric acid and potassium hydroxide are 
to be mixed in the proper proportions. The separate solu¬ 
tions contain respectively hydrogen and chlorine ions and 
potassium and hydroxyl ions. When the solutions are 
mixed, the hydrogen and hydroxyl ions immediately unite 
to form molecules of water, because water does not dis¬ 
sociate into ions to any appreciable extent. The final solu¬ 
tion is thus rendered neutral by the removal of the hydrogen 
and the hydroxyl ions — the acidic and basic constituents 
respectively. 


134 


A BRIEF COURSE IN CHEMISTRY 


The ionic equation expressing the neutralization of potas¬ 
sium hydroxide by hydrochloric acid is: — 

K+ + OH- + H+ + Cl- = K+ + Cl- + H 2 O 

The potassium and chlorine ions move freely about in the 
solution. If the solution is evaporated, the ions gradually 
unite as the solution becomes more and more concentrated, 
until finally nothing remains except the neutral salt, solid 
potassium chloride. Since neutralization is the combining 
of hydrogen and hydroxyl ions to form water, the general 
ionic equation for neutralization is: — 

H+ + OH- = H 2 O 

Hydrogen Ion Hydroxyl Ion Water 

176, Interpretation of certain facts by the ionic theory. — 

The ionic theory, like other theories, must meet one im¬ 
portant requirement, f.e., it must explain facts derived from 
experiment. We shall now interpret certain facts by the 
ionic theory. 

177. Behavior of solutions toward an electric current. — 

Water itself conducts electricity very slightly indeed. If 
solutions are subjected to the action of an electric current, 
the results vary. Solutions of some substances conduct 
electricity readily; these substances are acids, bases, and 
salts. Whereas solutions of other substances do not conduct 
electricity at all. 

A simple experiment enables us to find out what substances are 
electrolytes and what are non-electrolytes, i.e., what substances form 
conducting solutions and what form non-conducting. We have already 
learned that solutions of acids, bases, and salts conduct electricity 
(§ 164). We can proceed in the same way with other soluble substances. 
The apparatus is the same as that shown in Fig. 69. The procedure is 
the same, except that solutions of sugar, alcohol, glycerin, etc., are put 
in the vessel A. The lamp will not glow, showing that the only solu¬ 
tions that conduct electricity are those from acids, bases, and salts. 

This fact is readily interpreted by the ionic theory. Solu¬ 
tions of acids, bases, and salts contain electrically charged 
particles — ions. Whereas solutions of other substances do 
not. Hence, when an electric current is introduced into 


IONS AND IONIZATION 


135 


solutions of acids, bases, or salts, particles already charged 
with electricity are there to conduct the current. No such 
particles are in solutions of other substances; hence such 
solutions can not conduct an electric current. 

178. Ionic interpretation of the chemical reactions in 
solutions of acids, bases, and salts. — Dry potassium 
chloride and dry silver nitrate do not interact chemically, 
but if their solutions are mixed, a precipitate of silver chlo¬ 
ride is immediately produced. Furthermore, any dissolved 
chloride will interact in the same way with silver nitrate or 
any soluble silver salt. 

Let us interpret this fact by the ionic theory. Solutions 
of potassium chloride and silver nitrate contain potassium 
ions (K+), chlorine ions (Cl~), silver ions (Ag+), and nitrate 
ions (N03 ~). Now when certain pairs of ions are introduced 
into the same solution, they react chemically. Thus, silver 
ions and chlorine ions combine to form insoluble silver 
chloride; and this precipitate serves as evidence of the 
chemical change. According to the ionic theory, the source 
of the silver ions and the chlorine ions is immaterial. And 
it is a fact that hydrochloric acid and solutions of different 
chlorides precipitate silver chloride from any solution con¬ 
taining silver ions, e.g. , silver nitrate or silver sulfate. This 
is the explanation offered by the ionic theory for the pre¬ 
cipitation of silver chloride, which you will recall is the usual 
test for hydrochloric acid and all soluble chlorides (§ 136). 
The test is really a test for ions (Cl” and Ag+), not for sub¬ 
stances. 

The ordinary and ionic equations for this test are — 

HCl + AgNOg = AgCl + HNO 3 
H+ + Cl- + Ag+ + NO 3 - = AgCl + H+ + NO 3 - 
Cl- + Ag+ = AgCl 

No chemical action is observed when solutions of potassium 
chlorate and silver nitrate are mixed, because potassium 
chlorate solution contains potassium ions (K+) and chlorate 
ions (CIO3-) — not chlorine ions (CD). Hence when silver 
nitrate solution is added, no precipitate is formed because 


136 


A BRIEF COURSE IN CHEMISTRY 


silver nitrate is effective in testing for only those chlorine 
compounds that yield chlorine ions, not for other chlorine 
compounds, such as potassium chlorate (KCIO 3 ). 

Similar to the chloride test, the test for sulfuric acid and 
all soluble sulfates is the formation of insoluble barium sul¬ 
fate (BaS 04 ) upon the addition of a solution of barium chlo¬ 
ride or any other soluble barium compound. Sulfuric acid 
and sulfate solutions contain sulfate ions (SO4 ), which 
combine with barium ions (Ba++) furnished by the soluble 
barium compound. But this test is not applicable to 
other sulfur compounds, such as sulfides, sulfites, and thio¬ 
sulfates, because solutions of these sulfur compounds do 
not contain sulfate ions. The ionic equation for a sulfate 
test is: — 


Ba++ + 2 Cl- + 2 H+ + SO 4 -- = BaS 04 + 2 01“ + 2 H+ 

179. Other properties besides the formation of precipitates 

are due to ions. — Let us consider two cases. (1) The sour 
taste of all acids is attributed to hydrogen ions (H+), which 
are common to acids. (2) The color of certain solutions is 
also due to ions. Most ions are colorless, whereas solutions 
having a common colored ion have the same color. Thus, 

copper ions (Cu++) are blue, 
and solutions of copper salts 
are blue, irrespective of the 
color of the undissolved cop¬ 
per compound. Similarly, 
cobalt ions (Co++) give pink 
and nickel ions (NU*^) green 
solutions. 

180. Electrolysis and ions. 

— The simple experiments 
with solutions of acids, bases, 
and salts described in § 164 
(in which the bulb glowed) are examples of electrolysis. 
Electrolysis is the term applied to the series of changes ac¬ 
companying the passage of an electric current through a 
solution of an acid, base, or salt. 

If an electric current is passed through a solution of hydro- 



Fig. 71. — Sketch to illustrate the 
electrolysis of hydrochloric acid. 






















IONS AND IONIZATION 


137 


chloric acid, bubbles of gas rise from each electrode — 
hydrogen from the cathode or negative electrode and chlorine 
from the anode or positive electrode (Fig. 71). 

Let us interpret this example of electrolysis by the ionic 
theory. 

(1) The solution of hydrochloric acid contains some un¬ 
dissociated hydrochloric acid molecules (HCl), positively 
charged hydrogen ions (H"*"), and negatively charged chlo¬ 
rine ions (Cl“). These ions are in the solution before the 
electric current is passed through the solution. 

(2) When the electric current is passed in, the electrodes 
become charged with electricity — the anode positively 
(+) and the cathode negatively (—). 

(3) According to an old principle, bodies charged with the 
same kind (e.g., plus) of electricity repel each other; whereas 
bodies charged with different kinds (Lc., plus and minus) 
attract each other. In other words, the negative electrode 
(cathode) attracts the positive ions (cations) and repels 
the negative ions (anions). At the same time the positive 
electrode (anode) attracts the negative ions (anions) and 
repels the positive ions (cations). As a result of this attrac¬ 
tion and repulsion the ions actually move toward the elec¬ 
trodes and, of course, carry their electric charges with them. 
In brief, anions go to the anode and cations to the cathode. 

In this case, hydrogen ions (H+) go to the cathode and 
chlorine ions (Cl“) go to the anode (Fig. 71). This migration 
of ions, as it is called, was first studied carefully by the 
English scientist Faraday (Fig. 72), who named the charged 
particle ion ” from a Greek word meaning wanderer.’’ 
Thus, the term ion emphasizes the fact that when the current 
of electricity is passed through the solution, the two kinds of 
electrically charged particles (already in the solution) move 
toward their electrodes in two more or less orderly proces¬ 
sions — negative ions migrating to the positive electrode 
and positive ions to the negative electrode. 

(4) As soon as the ions come in contact with their respec¬ 
tive electrodes they act in accordance with another long- 
established principle; ms., they lose their electric charges. 
In other words, when the positive ions, or cations (H+), 




138 A BRIEF COURSE IN CHEMISTRY 

reach the cathode, the electric charges on the ions are neu¬ 
tralized. Electric charges, equal in quantity, though oppo¬ 
site in kind, are lost by the cathode, but are constantly re¬ 
newed by the battery or dynamo. The hydrogen ions (H''") 
once deprived of their electric charges do not regain them, 
but immediately become ordinary, uncharged hydrogen 
atoms (H) which combine and escape as molecules of hydro¬ 
gen (H 2 ). Simultaneously, the negative ions, or anions 
(Cl“), migrate to the anode, lose their charges, become chlo¬ 
rine atoms (Cl), unite, 
and escape as chlorine 
molecules (CI 2 ). As a 
final result of the elec¬ 
trolysis of hydrochloric 
acid, hydrogen and chlo¬ 
rine are the sole products. 

181. Electrolysis often 
yields secondary prod¬ 
ucts. — Let us consider 
two cases. 

1. When a solution of 
sodium hydroxide under¬ 
goes electrolysis, hydro¬ 
gen is liberated at the 
cathode and oxygen at 
the anode. The electrol¬ 
ysis proceeds as usual, 
the cations (sodium ions, 
Na+) migrate to the 
cathode and the anions 
(hydroxyl ions, OH~) to the anode. But when the ions are 
discharged, chemical changes take place. The sodium atoms 
(Na), which are produced at the cathode, react with the water 
(in the solution), forming hydrogen and sodium hydroxide; 
the hydrogen escapes and the sodium hydroxide remains — 
dissolved in the solution. The unstable hydroxyl radicals 
(OH), which are produced at the anode, break down into 
oxygen (O 2 ) and water (H 2 O). Hence, the final result is the 
production of two secondary products; hydrogen at the 






IONS AND IONIZATION 


139 


cathode and oxygen at the anode. We might summarize 
the process as follows: — 

( 1 ) Ionization, NaOH = Na+ + OH~ 

( 2 ) At the cathode, 2 Na + 2 H 2 O = H 2 + 2 NaOH 

(3) At the anode, 4 OH = 02 + 2 H 2 O 

2. In a copper sulfate solution the ions are copper ions 
(Cu++) and suKate ions (SO4 ). When the electric current 
is passed through the solution, the copper ions (Cu++) migrate 
to the cathode, lose their electric charges, become copper 
atoms (Cu), and adhere as metallic copper to the cathode. 
The sulfate ions (SO4 ) migrate to the anode, lose their 
electric charges, and immediately interact with the water 
around the anode, forming sulfuric acid (H2SO4), which 
remains in the solution (as ions), and oxygen atoms (O), 
which unite into molecules (O2) and escape from the solution. 

The so-called electrolysis of water (§ 97) is, strictly speaking, not the 
actual electrolysis of water itself but of a solution of sulfuric acid. It 
is interpreted as follows: Water itself does not conduct electricity, but 
if sulfuric acid is added, the solution is provided with hydrogen ions 
(H+) and sulfate ions (SO 4 ), which behave as just described. After 

the electrolysis, the solution contains the same weight of sulfuric acid 
but less water — in fact, the loss of water equals the weight of the liber¬ 
ated gases. 

182. Another definition of electrolysis. — We may now 

re-define or describe electrolysis as ionic migration induced 
by an electric current, the ions moving to their respective 
electrodes, where they are transformed into atoms or radicals 
which escape wholly or in part as elements, or which form 
various products by interaction with the water of the solution. 

183. What ions are in a solution? — As a rule the ions 
formed by an electrolyte are the charged atom, or group, 
corresponding to the two parts of the compound. Thus, 
nitric acid consists of the two parts H and NO 3 , and the ions 
are H+ and NOs”. Similarly, the parts of sodium hydroxide 
are Na and OH, and the ions are Na+ and 0H“. In most 
cases the name of the compound is made up of the parts, 
e.g.j copper sulfate, calcium hydroxide. 


140 


A BRIEF COURSE IN CHEMISTRY 


The ions normally formed by the ionization of acids, bases, 
and salts are given in Table IV. 

TABLE IV.— Ions 


Element 

OR 

Radical 

Ion 

Element 

OR 

Radical 

Ion 

Element 

OR 

Radical 

Ion 

Ammonium 

NH4+ 

Barium 

Ba++ 

Aluminum 

A1+++ 

Hydrogen 

H+ 

Calcium 

Ca++ 

Antimony 

Sb+++ 

Mercury (ous) 

Hg+ 

Copper 

Cu++ 

Bismuth 

Bi+++ 

Potassium 

K+ 

Iron (ous) 

Fe++ 

Iron (ic) 

Fe+++ 

Silver 

Ag+ 

Lead 

Pb++ 



Sodium 

Na+ 

Magnesium 

Mg++ 

Tin (ic) 

Sn++++- 



Mercury (ic) 

Hg++ 



Bromide 

Br- 

Tin (ous) 

Sn++ 



Chlorate 

CIO3- 

Zinc 

Zn++ 



Chloride 

ci- 





Hydroxyl 

OH- 

Carbonate 

CO3-- 



Iodide 

I~ 

Chromate 

Cr04-- 



Nitrate 

NO3- 

Dichromate 

CraOr-- 



Sulfate (acid) 

HSO4- 

Sulfate 

SO4-- 





Sulfide 

S-- 





Sulfite 

S03-- 




184. What electric charges are on ions ? — As repeatedly 
stated, cations are positive and anions are negative. These 
can be selected, of course, from the table of ions. It is con¬ 
venient, however, to remember these simple rules regarding 
the kind of charge : — 

(1) Hydrogen and metals form positive ions or cations 
{e.g.y H+, Cu++). Ammonium forms a positive ion (NH4+). 

(2) Hydroxyl forms a negative ion (OH“). 

(3) Non-metals — except hydrogen — form negative ions 
or anions {e.g., Cl~). 

(4) Characteristic radicals of acids and their corresponding 
salts form negative ions or anions {e.g., NOa", S 04 ~~). 

The number and kind of electric charges on an ion are the 
same as the valence of the corresponding atom or radical 
(§ 117). Thus Na+ and Na+h Ba++ and Ba+^ 804““ and 

















IONS AND IONIZATION 


141 


S 04 “^ AI+'++ and Al+^. Moreover, just as the total valences 
of the two parts of a compound are equal, so the total number 
of each kind of charges (+ and —) is equal. For example, 
Na+i CH and Na+, Cl~; Al 2 +^+^ ( 804 ) 3 “^“^“^ and 2 A1+++, 
3 S 04 ~ (Note : The coefficient multiplies the sign as well 
as the symbol, e.g., 2 A1+++ = 2 A1 and + 6.) Valence is 
defined as the number of electric charges carried by an ion. 

185. Degree of ionization of acids, bases, and salts. — 
Except in unusual cases solutions of acids, bases, and salts 
contain both molecules and ions. That is, not all the mole¬ 
cules dissociate into ions. The degree of dissociation depends 
upon the concentration of the solution and also upon the 
electrolyte itself. In concentrated solutions the number of 
molecules is large. As the solution is diluted, more and more 
molecules dissociate into ions. 

Strong acids and bases are almost completely dissociated 
in dilute solutions, and weak ones are only slightly disso¬ 
ciated. We do not call salts “ weak or “ strong, because 
most of them dissociate to about the same degree. 

The approximate per cent of dissociation of certain acids, 
bases, and salts in rather dilute solutions of the same relative 
concentration and at the same temperature is shown in 
Table V. 


TABLE V. — Per Cent of Ionization 


Substance 

Ions 

Per Cent of 
Ionization 

Hydrochloric Acid. 

H+ Cl- 

92 

Nitric Acid. 

H+ NO 3 - 

92 

Sulfuric Acid. 

2H+ SO 4 -- 

61 

Acetic Acid. 

H+ C 2 H 3 O 2 - 

1.3 

Sodium Hydroxide. 

Na+, OH- 

91 

Ammonium Hydroxide .... 

NH4+ OH- 

1.3 

Sodium Chloride. 

Na+ Cl- 

84 

Silver Nitrate. 

Ag+ NO 3 - 

81 

Barium Chloride. 

Ba++, 2 Cl- 

76 
















142 


A BRIEF COURSE IN CHEMISTRY 


EXERCISES 

1. State the ionic theory, and several facts that support it. 

2. Define and illustrate ion, anion, cation, electrode, anode, cathode, 
positive electrode, negative electrode. 

3. What is the difference between an atom and an ion of sodium? 
How is each represented ? 

4. Name the ions in a solution of (a) hydrochloric acid, (6) sodium 
chlorate, (c) calcium hydroxide, (d) sodium nitrate, (e) zinc sulfate. 

6 . State the chemical expression for these ionized elements or radi¬ 
cals : (a) Hydrogen, sodium, potassium, silver, ammonium; (6) chlo¬ 
ride, nitrate, hydroxyl; (c) calcium, barium, copper, zinc, magnesium, 
lead, iron (ous); (d) sulfate, carbonate; (e) aluminum, iron (ic). 

6 . Interpret neutrahzation by the ionic theory. Write the funda¬ 
mental equation for neutralization. 

7. Interpret by the ionic theory (a) test for a chloride ; (6) test for 
a sulfate. 

8 . Define acid, base, and salt in terms of the ionic theory. 

9. Define electrolysis. Interpret by the ionic theory the elec¬ 
trolysis of (a) hydrochloric acid, (b) sodium hydroxide, (c) copper 
sulfate. 

10. Write ionic equations for (a) potassium chloride and silver ni¬ 
trate form silver chloride and potassium nitrate, and (6) barium nitrate 
and sodium sulfate form barium sulfate and sodium nitrate. 


PROBLEMS (Review) 

(See also Problems at the end of Chapters VIII and IX.) 

1. If 36 gm. of copper are heated in air until there is no farther in¬ 
crease in weight, what is the name, formula, and weight of the product ? 
Ans. 45.07 gm. 

2. Equal weights of sodium and calcium interact with water, and 
the liberated gas is collected. Which metal yields the larger volume? 
Write the equation for each reaction. 

3. How many grams of zinc must be used with hydrochloric acid 
to produce 750 cc. of hydrogen at 0° C. and 760 mm. ? 


SUGGESTIONS FOR LABORATORY WORK 

♦ 

(References are to Newell’s Laboratory Exercises in Chemistry) 

Exercise 32 — Electrolytes and Non-Electrolytes — T. 

Exercise 33 — Electrolysis of Copper Sulfate Solution — T. 

Exercise S31 — Electrolysis of Copper Sulfate Solution (Short 
Method) — T. 


IONS AND IONIZATION 


143 


Exercise S32 — Electrolysis of Sodium Sulfate Solution — 
Exercise 34 — Reversible Reactions. 

Exercise 35 — Colored and Colorless Ions. 

Exercise S33 — Testing for Ions. 

Exercise S34 — Hydrolysis of Certain Salts. 

Exercise S56 — Tests for Metals. 

Exercise S59 — Testing Salts for Metal and Non-Metal. 


T. 


CHAPTER XV 


AMMONIA — AMMONIUM HYDROXIDE — 
AMMONIUM COMPOUNDS 

186. Introduction. — The term ammonia includes both 
the gas (NHs) and its solution (NH4OH). Sometimes the 
solution is called ammonia water, though its scientific name 
is ammonium hydroxide. 

187. Formation of ammonia. — Ammonia is liberated 
during the decomposition of nitrogenous matter, ^.e., vege¬ 
table and animal matter containing nitrogen. If hair, 
leather, fur, and feathers are heated (especially with lime 
or soda-lime), ammonia is given off and can be detected with 
moist red litmus paper in the pungent gases. The formation 
of ammonia in this way is a test for combined nitrogen 
(§ 191, last paragraph). Soft coal contains combined nitro¬ 
gen and hydrogen, and when the coal is heated, as in making 
illuminating gas, ammonia is liberated. This is one source 
of ammonium hydroxide (§ 250). 

188. Preparation of ammonia in the laboratory. — Am¬ 
monia is prepared in the laboratory by heating ammonium 
chloride with a non-volatile hydroxide, such as moist calcium 
hydroxide. The equation is : — 

2 NH4CI + Ca(OH)2 = 2 NH4OH + CaCh 

Ammonium Calcium Ammonium Calcium 

Chloride Hydroxide Hydroxide Chloride 

The ammonium hydroxide is unstable, especially when 
heated, and decomposes into ammonia and water, thus: — 

NH4OH = NH3 + H2O 

Ammonium Hydroxide Ammonia Water 

The gas is very volatile, and is usually collected by allowing 
it to fiow upward into a bottle and displace the air. 

144 


AMMONIA —AMMONIUM HYDROXIDE 


145 


189. Manufacture of ammonium hydroxide. — Ammo¬ 
nium hydroxide is manufactured by passing ammonia into 
water. Much of the ammonia from which commercial 
ammonium hydroxide is manufactured is obtained by the 
distillation of coal, i.e., by heating the coal in closed retorts, 
as in the manufacture of illuminating gas and coke (§§ 260, 
44). The ammonia is separated from the other gases by 
washing (or “ scrubbing ”) it out with water (§ 250 and 
Fig. 95). This impure solution, called ammoniacal liquor or 
gas liquor, is treated with lime, and the liberated ammonia is 
passed into tanks containing hydrochlo¬ 
ric acid or sulfuric acid. Lime is added 
to this solution of ammonium chloride 
or sulfate, and the liberated ammonia 
is led into water, forming thereby the 
ammonium hydroxide of commerce. 

Ammonium hydroxide is manufac¬ 
tured from the ammonia prepared by 
the direct combination of nitrogen and 
hydrogen (§§ 192, 193). 

190. Physical properties of ammonia. 

— Ammonia is a colorless gas. It has 
an exceedingly pungent odor, and if 
inhaled suddenly or in large quantities, 
it brings tears to the eyes and may 
cause suffocation. It is light and vola¬ 
tile, being about one-half (0.59) as heavy 
as air. A liter of the gas at 0° C. and 
760 mm. weighs 0.77 gm. 

Ammonia gas is easily liquefied; the usual conditions are 
0° C. and 4.2 atmospheres (i.e.y 4 X 760 mm.). Liquid, or 
liquefied, ammonia is often called anhydrous ammonia, 
because it contains no water. It boils at — 34° C. Hence, 
if it is exposed to the air or warmed in any way, it changes 
into the gas, and in so doing absorbs considerable heat. 
This fact has led to the extensive use of liquid ammonia in 
refrigeration and in making ice (§ 194). 

Ammonia gas is very soluble in water. Its solubility can 
be shown by the fountain experiment. (See Fig. 73 and com- 



Fig. 73. — Experiment 
to show the marked 
solubility of ammonia 
in water. 





















146 


A BRIEF COURSE IN CHEMISTRY 


pare § 131.) A liter of water at ordinary temperature dis¬ 
solves about 700 1. of gas. 

A solution of ammonia gives off the gas freely, as may be 
easily discovered by the odor or by the formation of dense 
white fumes of ammonium chloride (NH 4 CI) when the solu¬ 
tion is exposed to hydrochloric acid (§ 132). 

191. Chemical properties of ammonia. — Ammonia will 
not burn in air under ordinary conditions, nor will it support 
combustion, as the term is usually used. 

Ammonia reacts with certain elements, e.g., magnesium 
and chlorine, thus: — 

2 NH 3 -f" 3 Mg = Mg3N2 “b 3 H 2 

Ammonia Magnesium Magnesium Nitride 

2NH3 + 3CI2 = N2 + 6 HC 1 

Ammonia Chlorine Nitrogen Hydrochloric Acid 

Ammonia combines directly with water, thus: — 

NH3 + H2O = NH4OH 

Ammonia Water Ammonium Hydroxide 

Ammonia unites with acids and forms ammonium salts, 
e.g., NH3 and HCl (gas) form ammonium chloride (NH4CI). 
This reaction serves as a test for ammonia (§ 132). 

192. Synthesis of ammonia from its elements. — Am¬ 
monia can be prepared by synthesis, i.e., by direct union of 
nitrogen and hydrogen. The two gases unite if electric 
sparks are passed through their mixture. The equation is: — 

N2 + 3 H2 = 2 NH3 

The amount of ammonia formed in a given case, however, 
is only a small per cent of that indicated by the equation. 
The small yield is due partly to the fact that ammonia itself 
decomposes into nitrogen and hydrogen, thus: — 

2 NH3 = N2 + 3 H2 

If we compare these equations, we see that one is the reverse 
of the other. When the experiment is done in a closed tube, 
the two reactions proceed at the same time — one reversing 
the other. Such a complete reaction is called a reversible 


AMMONIA — AMMONIUM HYDROXIDE 147 

reaction. The equation for a reversible reaction contains 
oppositely pointed arrows, thus: — 

N 2 + 3H2::^2NH3 

This equation is read: nitrogen and hydrogen react reversibly 
to form ammonia. 

A reversible reaction under a given set of conditions pro¬ 
ceeds to equilibrium. This means that the amounts of the 
substances involved in both reactions increase or decrease 
until the quantity of any one substance formed equals the 
quantity of it which is transformed in a given time. In 
the case of ammonia, the mixture at equilibrium is only 
about 2 per cent ammonia, the rest being nitrogen and 
hydrogen. 

If ammonia is removed by introducing acid or water into 
the apparatus, or by some other device, its removal dis¬ 
places the equilibrium (toward the NH 3 ) and the reaction pro¬ 
ceeds to completion, ^.e., all, or practically all, the nitrogen 
and hydrogen combine. 

193. Manufacture of ammonia from nitrogen and hydro¬ 
gen by S 5 mthesis. — The reaction described in § 192 for the 
manufacture of ammonia proceeds slowly. To be profitable 
commercially, the reaction must be hastened, that is, its 
velocity must be increased so that more ammonia will be 
formed in a given time. Several factors affect the velocity 
of a reaction, e.g., temperature, pressure, and the presence 
of a catalyst. In manufacturing ammonia, the best condi¬ 
tions are a temperature of about 400° C. and a pressure of 
about 200 atmospheres (f.e., 200 X 760 mm.). Under these 
conditions the yield is from 5 to 8 per cent. 

Even under these conditions the velocity of the reaction 
is too slow for industrial use. Hence a catalyst is used to 
speed up the reaction. One catalyst is iron containing small 
amounts of potassium and aluminum oxides. Ammonia 
made directly from its elements is called synthetic ammonia. 

During the World War large quantities of ammonia were 
manufactured in Germany by this process, called the Haber 
process. It is in operation in Germany and also in the 
United States at the present time. 


148 


A BRIEF COURSE IN CHEMISTRY 


A diagram of the essential parts of the apparatus for manufacturing 
ammonia from nitrogen and hydrogen is shown in Fig. 74. The mixture 
of purified nitrogen and hydrogen is passed through A into the com¬ 
pression vessel B, and forced through C into the reaction chamber H. 
Here the mixture comes in contact with the catalyst and is changed in 
part into ammonia. The mixture of ammonia and uncombined gases 
passes into the coiled pipe in the vessel E which contains a circulating, 
very cold liquid (entering at J and flowing out at K). The ammonia 
condenses to a hquid and accumulates in F, while the uncombined 
nitrogen and hydrogen pass along through the valve G and the pipe H 
into D again. The liquefied ammonia is drawn off through I. 

194. Liquefied ammonia as a refrigerant. — The use of 
liquefied ammonia in producing low temperatures in cold 



Fig. 74. — Diagram of apparatus for manufacture of ammonia from nitrogen 
and hydrogen. 

storage plants and refrigerators (see § 221 ) depends upon 
two facts. (1) Liquefied ammonia (not ordinary ammonia 
solution) changes rapidly into a gas when its pressure is 
reduced, and (2) in so doing absorbs heat from the surround¬ 
ing air or liquid. Hence, if liquefied ammonia is allowed to 
flow through a pipe immersed in a solution of sodium chloride 
or calcium chloride (technically called a brine), the ammonia 
evaporates in the pipe and cools the brine, which may be 
used directly as a refrigerant or for making ice. In some 
cold storage plants, packing houses, and sugar refineries, this 
cold brine is circulated through pipes placed in the storage 
rooms where a low temperature is desired (Fig. 75). 

The construction and general operation of an ice-making plant is 
shown in Fig. 76. Liquefied ammonia is forced from a tank into a 
series of pipes which are submerged in a large vat A nearly filled with 
brine. Metal cans containing purified water to be frozen are immersed 



























AMMONIA — AMMONIUM HYDROXIDE 


149 


in the brine, which is kept below the freezing point of water by rapid 
evaporation of the ammonia in the pipes. After several hours the water 
in the cans is frozen into cakes of ice. As fast as the ammonia gas forms 



Fig. 75. — A room in a cold storage plant. 


in the pipes, it is removed by exhaust pumps E into another set of 
pipes C, where it is condensed into liquefied ammonia and conducted 
through D into the other set of pipes ready for renewed use. 

In cold storage plants the cold brine is circulated through pipes F to 
the various rooms B. 


196. Ammonium hydroxide. — Ammonia combines with 
water to some extent and forms a solution of ammonium 



Fig. 76. — Apparatus for using liquefied ammonia to produce low tempera¬ 
tures. 




































































150 


A BRIEF COURSE IN CHEMISTRY 


hydroxide (NH 4 OH). The NH 4 acts as a radical and in a 
solution becomes NH 4 '^ just as K in KOH becomes K+. 

Ammonium hydroxide is a base (§ 161), though a much 
weaker one than sodium hydroxide (§ 185), in spite of its 
pungent odor. Like other bases it turns litmus blue. Con¬ 
centrated solutions have a slippery feeling. It also neutral¬ 
izes acids, thus: — 

NH 4 OH + HCl = NH 4 CI -h H 2 O 

Ammonium Hydrochloric Ammonium Water 

Hydroxide Acid Chloride 

Ammonium hydroxide is widely used as a cleansing agent 
(especially for the removal of grease), and is sold under the 
names “ ammonia ” and ‘‘ household ammonia.’^ Large 
quantities are consumed in the manufacture of dyestuffs, 
sodium bicarbonate (§ 279), and ammonium compounds. 

196. Ammonium salts. — These salts, like the base am¬ 
monium hydroxide, contain a group of atoms which acts 
chemically like an atom .of a metal. This group is called 
ammonium, and its formula is NH 4 ; in solution it is (NH 4 )’''. 
Ammonium has not been isolated from its compounds. 
Ammonium, like hydroxyl (§ 161, last paragraph), is a radical, 
because it is the root or foundation of a series of compounds 
and in many chemical changes passes as a unit from one 
ammonium compound to another. 

Ammonium salts decompose when heated with sodium 
hydroxide or moist calcium hydroxide (§ 188) ; ammonia is 
the conspicuous product. This reaction is a test for ammo¬ 
nium compounds. Ammonium salts are highly ionized in 
solution and give the ion NH 4 '^. 

There are many ammonium salts. Ammonium chloride (NH4CI) is 
formed by the neutrahzation of ammonium hydroxide (a base) by hydro¬ 
chloric acid. It is manufactured by passing ammonia into hydrochloric 
acid. The crude product is often called muriate of ammonia to indicate 
its relation to muriatic acid (the commercial name of hydrochloric acid). 
It is used in some types of batteries {e.g., dry cells), and as an ingredient 
of soldering fluids and of fertilizers. 

Crude ammonium chloride is purified by heating it gently in a large 
iron pot with a dome-shaped cover; the ammonium chloride volatilizes 
and then crystalhzes quickly in the pure state as a fibrous mass on the 
inside of the cover; the non-volatile impurities remain behind in the 



AMMONIA — AMMONIUM HYDROXIDE 


151 


vessel. This process of purification is called sublimation. The product 
is a sublimate. Sublimed ammonium chloride is often called sal am¬ 
moniac. 

Ammonium sulfate ((NH 4 ) 2 S 04 ) is a grayish or yellowish solid. It 
is used as an ingredient of fertilizers, since it is a cheap, soluble salt con¬ 
taining considerable nitrogen (§ 143). 

EXERCISES 

% 

1. How is ammonium hydroxide manufactured? 

2. State the conspicuous properties of ammonia. 

3. Define and illustrate (by ammonia) (a) reversible reaction, 
(6) equilibrium, (c) displacement of equilibrium, (d) catalyst. 

4. State these reactions in the form of equations: (a) preparation of 
ammonium hydroxide from ammonium chloride and calcium hydroxide, 
(6) decomposition of ammonium hydroxide into ammonia and water, 
(c) nitrogen and hydrogen react reversibly to form ammonia. 

6. Describe the manufacture of ice by liquid ammonia. 

6 . State the test for (a) ammonia and (6) ammonium compounds. 

7. Complete and balance: (a) NH4OH -|-= (NH 4 ) 2 S 04 + 

-; (6) NH 3 +- = N 2 + HCl. 


PROBLEMS 

1. How many grams of ammonia (NH 3 ) can be obtained from 1 kg. 
of ammonium chloride and sufficient calcium hydroxide? 

2. What weight of ammonium chloride (95 per cent pure) is needed 
for the preparation of 60 gm. of NH 3 ? 

3. How many grams of ammonium chloride can be made from 
ammonium hydroxide and 100 gm. of the necessary acid? 

4. What formula corresponds to the composition N = 26.17 per cent, 
H = 7.48, Cl = 66.35? 

SUGGESTIONS FOR LABORATORY WORK 

(References are to Newell’s Laboratory Exercises in Chemistry) 

Exercise *36 — Preparation and Properties of Ammonia. 

Exercise S35 — Preparation of Ammonia from Various Substances, 


CHAPTER XVI 


NITRIC ACID — NITRATES 

197. Preparation of nitric acid. — Nitric acid is prepared 
in the laboratory by heating concentrated sulfuric acid with 
sodium nitrate (NaNOs). 

The chemical change at a moderate temperature is ex¬ 
pressed by the equation : — 

NaNOs + H2SO4 = HNO3 + NaHS04 

Sodium Sulfuric Nitric Acid Sodium 

Nitrate Acid Acid Sulfate 

At a high temperature and with an excess of sodium nitrate, 
the equation is: — 

2 NaNOs + H2SO4 = 2 HNO3 + Na 2 S 04 

198. Manufacture of nitric acid. — Nitric acid is manu¬ 
factured by several processes. 

1. The first is much like that used in the laboratory. A 
sketch of the apparatus is shown in Fig. 77. 

Sulfuric acid and sodium nitrate are heated in the cast iron retort 
(at the left) which is connected with quartz or stoneware tubes in which 
the vapor is condensed by cooling the tubes with a current of cold water 
(middle of sketch). The tubes are arranged so that the nitric acid first 
condensed runs into a reservoir for receiving concentrated acid. The 
vapors pass up an absorbing tower (at the right) where they are dissolved 
by descending water and flow out at the bottom as dilute acid. 

2. A second method uses air. It depends on the fact that 
nitrogen and oxygen unite if subjected to the temperature of 
an electric arc. Three main reactions are involved : — 

(1) N 2 + O 2 = 2NO 

Nitrogen Oxygen Nitric Oxide 

152 


NITRIC ACID —NITRATES 


153 


(2) 2 NO + O 2 = 2 NO 2 

Nitric Oxide Oxygen Nitrogen Dioxide 

(3) 3 NO 2 + H 2 O = 2 HNO 3 + NO 

Nitrogen Dioxide Water Nitric Acid Nitric Oxide 

Two unusual conditions must be fulfilled in this process. 
First, the nitrogen and oxygen must be heated to a very 
high temperature (about 3000° C.) before they will unite to 
an appreciable extent. Second, the mixture of gases result- 



Fig. 77. — Sketch of the apparatus for manufacturing nitric acid from 
sodium nitrate and sulfuric acid. 


ing from reaction (I) must be cooled very quickly. Reaction 
(I) is reversible, thus: — 

N 2 “h O 2 ^ 2 NO 

The maximum quantity (only about 5 per cent) of nitric 
oxide is obtained at about 3000° C. At this temperature 
equilibrium is reached, i.e., the relative proportions remain 
unchanged. (Compare § 192.) But just as soon as the 
temperature becomes lower, the reverse reaction (right to 
left) begins and the quantity of nitric oxide rapidly decreases. 
However, if the mixture is cooled very quickly, enough nitric 
oxide is left for reaction (2). 

Reaction (2) takes place readily; the nitric oxide (NO) 
liberated from (1) combines with more oxygen to form nitro¬ 
gen dioxide (NO 2 ), so that little nitric oxide (NO) is lost. 

A sketch of one form of the apparatus is shown in Fig. 78 and certain 
parts are shown in detail in Fig. 79. Air is blown (by A) into the elec- 

















































154 


A BRIEF COURSE IN CHEMISTRY 


trie furnace B. Here it is raised to the proper temperature (about 
3000° C.) by passing through an electric arc. In order to provide a 
large heating surface, the arc is spread out by powerful electromagnets 



Fig. 78. — Sketch of the apparatus for the manufacture of nitric acid from 
nitrogen, oxygen, and water. 


into a disk six feet in diameter (Fig. 79, left). The electrodes usually 
are hollow and are kept cool by circulating water; graphite electrodes 
are also used. An end view (Fig. 79, right) shows the magnets and edge 
of the disk. The hot gases containing the nitric oxide (NO) from the 
furnace are suddenly cooled in C, pass through the boiler D into the 
oxidizing chamber E, where nitrogen dioxide (NO 2 ) is formed. 

The nitrogen dioxide passes from E into the tower F, which is filled 
Mth tiles over which water trickles. Here reaction (3) takes place and 
dilute nitric acid is produced. The dilute acid is concentrated or con¬ 
verted (by limestone or lime) into calcium nitrate. The latter is used 
as a fertilizer, either alone or mixed with lime. 




Fig. 79. — Electric arc — spread out (left) and end view (right). 


3. In another process, which is rapidly coming into use, 
ammonia is the starting point. A heated mixture of ammonia 
and air is passed into fire-brick-lined chambers containing a 























































NITRIC ACID —NITRATES 155 

catalyst — usually platinum in the form of gauze. At about 
1000° C., the reaction is: — 

4 NHs + 5 O 2 = 4 NO + 6 H 2 O 

The nitric oxide is cooled, mixed with air, and then, as in 
the other process, passed into towers through which water 
trickles, where the nitric acid is produced, thus: — 

4 NO + 3 O 2 + 2 H 2 O = 4 HNO 3 

199. Physical properties of nitric acid. — Pure nitric acid 
is a colorless liquid, but the commercial acid is often yellow 
or brownish. So also, the acid that has been exposed to 
the sunlight is often yellow or brown, and if the light is in¬ 
tense, a brownish gas may often be seen in the bottle. It is 
somewhat volatile, and the vapor dissolves readily in water; 
hence the acid forms irritating fumes when exposed to air, 
especially moist air. (Compare § 311.) 

Commercial concentrated nitric acid contains about 68 
per cent of the compound HNO 3 , the rest being water. 

200. Chemical properties of nitric acid. — Nitric acid is 
sour, turns blue litmus red, and forms salts — the ni¬ 
trates. It is an unstable compound, and decomposes 
readily; among the decomposition products is the gas nitro¬ 
gen dioxide (NO 2 ). 

Nitric acid is a very corrosive substance and reacts readily 
with many substances. With nitrogenous organic sub¬ 
stances, ^.e., organic substances containing nitrogen, like 
hair, feathers, wool, silk, finger nails, and skin it forms a 
yellow substance. This change is. sometimes used as a 
test for nitric acid (§ 203, last paragraph). The concentrated 
acid causes serious burns and if spilled on the hands or face 
should be washed off immediately. With certain organic 
compounds it forms nitro-derivatives, such as nitroglycerin 
and nitrocellulose. 

One of the decomposition products of nitric acid is oxygen. 
Hence nitric acid is an oxidizing agent. Hot charcoal burns 
brilliantly in the hot acid, while straw, sawdust, hair, and 
similar substances are charred and even set on fire by it. 


156 


A BRIEF COURSE IN CHEMISTRY 


Some organic compounds, when heated with nitric acid, are 
completely decomposed into carbon dioxide and water. 

In the mixture of concentrated nitric and hydrochloric 
acids called aqua regia, nitric acid acts as an oxidizing agent 
(§ 136). The equation is usually written: — 

HNO 3 + 3HC1 = 2 H 2 O + CI 2 + NOCl 

Nitric Hydrochloric Water Chlorine Nitrosyl 

Acid Acid Chloride 

Nitric acid interacts readily and often violently with 
metals, metallic oxides, and hydroxides. (Compare § 132.) 
The products of these reactions vary, the chief ones being 
nitrates and nitrogen oxides (§ 203). 

201. Uses of nitric acid. — Nitric acid is one of the com¬ 
mon laboratory acids. Large quantities are used in the 
manufacture of nitrates, plastic substances, dyestuffs, sul¬ 
furic acid, cellulose nitrates, and explosives, e.g., guncotton, 
picric acid, and TNT (tri-nitrotoluene). 

202. Nitrates. — Nitric acid forms salts called nitrates. 
They are prepared by the methods usually used for salts, 
i.e.f the interaction of nitric acid and metals or metallic 
oxides and the neutralization of hydroxides by nitric acid. 

Many nitrates are white solids; but those of copper, nickel, 
and cobalt are blue, green, and dark red respectively. The 
nitrates of most metals are soluble in water. Their solu¬ 
tions are frequently used in the laboratory. The solids 
behave in various ways when heated. Equations illustrating 
typical reactions are: — 

2 NaNOs = 2 NaN02 “b O 2 j 

Sodium Nitrate Sodium Nitrite Oxygen 

2 Cu(N 03)2 = 2 CuO + 4 NO 2 + O 2 

Copper Nitrate Copper Oxide Nitrogen Dioxide Oxygen 

Since many nitrates, when heated, give up oxygen, they 
are powerful oxidizing agents. Thus, when potassium 
nitrate (KNO 3 ) is dropped on hot charcoal, the charcoal 
bums vigorously. This kind of chemical action is called 
deflagration. Nitrates are ingredients of gunpowder and 
fireworks. 


NITRIC ACID — NITRATES 


157 


203. The interaction of nitric acid and metals. — This 
action is exceedingly vigorous. The products of the reaction 
vary with the metal, the concentration of the acid, and the 
temperature. Hydrogen is not liberated as a rule so that 
it can be collected, for it is oxidized at once to water, whereas 
the nitric acid is reduced to nitrogen compounds — usually 
to nitric oxide (NO) with dilute acid and to nitrogen dioxide 
(NO 2 ) with concentrated acid. 

The interaction of nitric acid and copper will serve as an 
example of the common reactions. When moderately dilute 
nitric acid (sp. gr. 1.2) is poured on copper, a reddish brown 
gas is given off, and the liquid turns blue, owing to dissolved 
copper nitrate. The equation for the reaction is : — 

8 HNO 3 + 3 Cu = 3 Cu(N 03)2 + 2 no + 4 H 2 O 

Nitric Copper Copper Nitric Water 

Acid Nitrate Oxide 

Nitric oxide is represented as a product of the interaction 
of nitric acid and copper. But if the reaction takes place 
in an open vessel, the nitric oxide, which is a colorless gas, 
combines with oxygen and forms the reddish brown nitrogen 
dioxide gas. The equation is: — 

2 NO + O2 = 2NO2 

Nitric Oxide Oxygen Nitrogen Dioxide 

Hence we often speak of nitrogen dioxide as a product of 
the interaction of nitric acid and metals like copper, though 
it is a secondary product. 

With concentrated nitric acid, copper reacts thus: — 
4 HNO 3 + Cu = Cu(N 03)2 + 2 NO 2 + 2 H 2 O 

Nitric Acid Copper Copper Nitrate Nitrogen Dioxide Water 

Non-metals of course do not produce nitrates. With 
sulfur the reaction is represented thus ((1) is dilute and 
(2) is concentrated acid): — 


( 1 ) 

( 2 ) 


2 HNO 3 -h S = H 2 SO 4 -h 2 NO 
6 HNO 3 -h S = H 2 SO 4 -h 6 NO 2 -h 2 H 2 O 


158 


A BRIEF COURSE IN CHEMISTRY 


The test for a nitrate (and of course for nitric acid) is not 
made, as customary, by producing a precipitate, because the 
nitrate ion (NOs”) does not form insoluble compounds with 
other ions. It is a color test, and is made as follows: Add 
to the solution of the nitrate or the nitric acid in a test tube 
an equal volume of ferrous sulfate solution (freshly prepared 
from clean ferrous sulfate and cold water). Mix well. In¬ 
cline the test tube, and pour concentrated sulfuric acid 
cautiously down the side of the test tube. A dark colored 
layer appears where the two liquids meet, owing to the for¬ 
mation of a brown unstable compound which has the com¬ 
position (approximately) 3 FeS 04.2 NO. 

EXERCISES 

1. Summarize the properties of nitric acid. 

2. What is the test for (a) nitric acid, (6) a nitrate, (c) nitric oxide ? 

3. Describe the interaction of nitric acid and copper and state the 
equations. 

4. Describe fully the manufacture of nitric acid from (a) sodium 
nitrate, (6) nitrogen oxides, and (c) ammonia. 

6. Complete and balance: (a) CuCOs H-= Cu(N 03)2 H- 

+-; (0 HNO3 + C = NO2 +-+ CO2. 


PROBLEMS 

(See Problems at the end of Chapters VIII and IX.) 

1. Calculate the percentage composition of (a) nitric acid, (b) po¬ 
tassium nitrate, (c) sodium nitrate. 

2. Calculate the formula and give the name of the compound cor¬ 
responding to (a) O = 76.19, H = 1.58, N = 22.22; (b) O = 47.52, 
K = 38.61, N = 13.86. 


SUGGESTIONS FOR LABORATORY WORK 

(References are to Newell’s Laboratory Exercises in Chemistry) 

Exercise 37 — Preparation of Nitric Acid — T. 

Exercise *38 — Properties of Nitric Acid. 

Exercise *39 — Nitric Oxide and Nitrogen Dioxide. 

Exercise *40 — Test for Nitric Acid and Nitrates. 

Exercise S23 — Aqua Regia — T. 


NITRIC ACID — NITRATES 


159 


Exercise S36 — Nitrous Oxide — T. 

Exercise S28 — General Properties of Acids (Nitric acid part only). 

SUPPLEMENTARY SECTIONS FROM PART II 

379. Nitrogen oxides. 

380. Nitrous oxide. 

381. Nitric oxide. 

382. Nitrogen dioxide. 


CHAPTER XVII 


SULFUR AND SULFIDES 

204. Introduction. — Sulfur is an ingredient of gun¬ 
powder. Considerable is used in the manufacture of fire¬ 
works. Large quantities are consumed in the rubber in¬ 
dustry, especially in making automobile tires. In the paper 
industry, sulfur is employed in making the sulfite needed 
to convert the wood into pulp. Sulfur is a constituent of 
mixtures for killing insect pests. These mixtures, e.g., 
lime-sulfur spray, liberate sulfur upon the injurious insect. 
Sulfur itself is also used as an insecticide, especially for 
killing Phylloxera — an insect which destroys grapevines. 

Large quantities of sulfur are used in making sulfur com¬ 
pounds, e.g., sulfur dioxide (SO2) and carbon disulfide (CS2). 

205. Occurrence. — Large deposits of free sulfur occur 
in volcanic regions, such as Japan and Mexico. Other 
deposits, as in Sicily, Louisiana, and Texas, were doubtless 
formed by the action of microorganisms on calcium sulfate. 

Sulfur compounds are abundant, e.g., lead sulfide (galena, 
PbS), zinc sulfide (sphalerite or zinc blende, ZnS), mercuric 
sulfide (cinnabar, HgS), copper sulfide (chalcocite, CU2S, 
and chalcopyrite, CuFeS2), iron sulfide (iron pyrites, FeS2), 
and calcium sulfate (gypsum, CaS04 • 2 H2O). 

Volcanic gases often contain sulfur dioxide (SO2), and 
the water of sulfur springs contains hydrogen sulfide (H2S). 
Sulfur is a constituent of certain organic compounds present 
in onions, horseradish, mustard, and eggs. Some varieties 
of coal and petroleum contain sulfur compounds. 

206. The United States sulfur industry. — Enough sulfur 
for all domestic and most foreign uses is obtained from the 
large deposits in southern United States, especially Texas. 

160 


SULFUR AND SULFIDES 


161 


The sulfur deposits are from 500 to 1000 feet deep. The 
sulfur is forced to the surface by an ingenious method devised 
by the American chemist Frasch. 







































































































162 


A BRIEF COURSE IN CHEMISTRY 


A hole is drilled through the overlying soil, sand, clay, and rock into 
the sulfur deposits below (Fig. 80, left). The well, as it is called, is 
equipped with a set of concentric pipes (Fig. 80, right). The pipes, 
inclosed in a casing, are driven down through the hole. The use of the 
pipes is clear from the diagram on the right. Through the two larger 
pipes water at a high temperature (170° C.) and under 100 pounds pres¬ 
sure is forced down to the bottom (Fig. 81). This water flows out into 
the porous rock and melts the sulfur (melting point about 114° C.), which 
collects in a pool at the bottom of the pipes. Through the smallest pipe 
(air inlet) hot air is forced down and forms a froth with the sulfur. This 
froth, owing to the pressure of the hot water and air, rises through the 
“ suKur pipe ” (right) and flows out at the surface into large vats from 



Fig. 81. — Hot water pipes and pumps. 


which it is pumped into large wooden bins, where it cools and solidifies. 
The wooden part is then removed, leaving the sulfur as a huge block. 

A single well often produces over 500 tons of sulfur a day. Some of 
the blocks contain as much as 100,000 tons of sulfur. When the whole 
block is cold, the wooden sides of the bin are removed, the sulfur is 
blasted into fragments, and loaded into cars by steam shovels (Fig. 82). 
The American sulfur industry is conducted on a prodigious scale, and 
yields large quantities of pure sulfur (about 99.5 per cent). 

207. Physical properties of sulfur. — Sulfur is a pale 
yellow, brittle solid, which sometimes has a faint odor. It 
is insoluble in water and most acids. Most varieties of 




SULFUR AND SULFIDES 


163 



sulfur dissolve readily in carbon disulfide (CS 2 ), and in 
sulfur chloride (Lc., sulfur monochloride,” S 2 CI 2 ). Sulfur 
does not conduct electricity at all well. Its conductivity 
is lower than that of practically any other solid substance; 
this fact is sometimes expressed by saying sulfur is a good 
insulator. Nor does it conduct heat well, its heat 
conductivity being but one half that of cork and one fourth 
that of ice. The specific gravity of the solid is about 2, i.e.j 
it is about twice as heavy as water. 

When heated slowly, sulfur melts at 114.5° C. to a thin, 
pale yellow liquid. As the temperature rises, the liquid 


Fig. 82. — Loading sulfur into cars by a steam shovel (left), partly removed 
blocks of sulfur (center), derricks for drilling machinery (right). 

darkens and thickens; at about 160° C. it is dark brown 
and viscous, at about 230° C. it is black and too thick to 
fiow from the vessel, while at about 445° C. it becomes thin 
again, boils, and turns into yellow sulfur vapor. 

Sulfur vapor, if formed on a large scale and then cooled, 
condenses to a fine powder called flowers of sulfur. Molten 
sulfur if poured in cylindrical molds forms sticks called roll 
sulfur or brimstone. Molecules of sulfur vapor at about 
445° C. contain eight atoms (Ss), while at about 1800° C. the 
molecules contain two atoms (S 2 ). At about 2000° C. 
molecules and atoms are identical, i.e.j at this high temper¬ 
ature sulfur is monatomic (S), 





164 


A BRIEF COURSE IN CHEMISTRY 




208. Chemical properties of sulfur. — Sulfur is an active 
element. When heated, it combines readily and directly 
with many elements, especially oxygen. Thus, it ignites 
readily in the air and burns with a pale blue flame, forming 

sulfur dioxide (SO 2 ); if burned in oxy¬ 
gen, a little sulfur trioxide (SO 3 ) is also 
formed. Moist sulfur is oxidized to 
sulfuric acid slowly by exposure to air. 
Thus: — 

Fig. 83. — Orthorhombic 2 S + 3 O 2 + 2 H 2 O = 2 H 2 SO 4 
crystals of sulfur. 

It combines directly with most metals, 
forming sulfides; the reaction is often accompanied with 
much heat and light, as in the case of zinc, copper, and iron 
(§8). It also combines directly with carbon to form carbon di¬ 
sulfide (§49), and with chlorine to form sulfur chloride (S 2 CI 2 ). 

209. Different modifications of solid sulfur. — There are 
two forms of solid sulfur. 

(1) One form is called orthorhombic, or sometimes rhombic, 
sulfur (Fig. 83). If sulfur is dissolved in carbon disulfide 
and the solution evaporated slowly, sulfur is 

deposited as small crystals. Well-formed 
crystals have eight sides and belong to the 
orthorhombic system (Fig. 83). Crystallized 
native sulfur also is orthorhombic. Roll sul¬ 
fur and flowers of sulfur are orthorhombic, 
though the crystals are often so interlaced 
or so small that their shape is obscured. 

(2) Another variety, called monoclinic, 
is obtained by letting melted sulfur cool 
slowly. If sulfur is melted in a crucible and 
the excess of liquid is poured off as soon as 
crystals shoot out from the walls near the 
surface, the interior of the crucible wilTbe 
found to be lined with long, dark yellow, 
shining needles (Fig. 84). They are mono¬ 
clinic crystals of sulfur. In a few days they turn dull and 
opaque, and in time change into small orthorhombic crystals. 
Monoclinic sulfur is stable only if kept above 96° C. 



Fig. 84. — Section 
of a crucible 
showing mono¬ 
clinic crystals of 
sulfur. 










SULFUR AND SULFIDES 


165 


These two modifications of solid sulfur have different 
physical properties. Orthorhombic sulfur has the specific 
gravity 2.06 and melts at 112.8° C. (if heated rapidly). 
The corresponding values of monoclinic sulfur are 1.96 and 
119.25° C. 

210. Different forms of liquid sulfur. — If sulfur is heated 
above the viscous stage or boiled and then cooled quickly 
by pouring it into water, a tough, rubberlike, amber-colored 
solid is formed. It is called amorphous or plastic sulfur. It 
is insoluble in carbon disulfide. Plastic sulfur when first 
formed is non-crystalline and for this reason is sometimes 
called amorphous (“ without crystal form It is unstable, 
and soon becomes hard, brittle, and yellow; after consider¬ 
able time it changes in part into orthorhombic sulfur. 

211. Sulfur is an allotropic element. — The different modi¬ 
fications of sulfur are elementary sulfur, though they have 
different properties. Each burns to sulfur dioxide, and the 
same weight of each yields the same weight of sulfur dioxide 
(e.gr., 32 gm. of each yields 64 gm. of sulfur dioxide). Sulfur 
is an allotropic element (§ 52). 

HYDROGEN SULFIDE AND OTHER SULFIDES 

212. Occurrence of hydrogen sulfide. — This is the 
gaseous compound of sulfur and hydrogen with a notoriously 
bad smell. It occurs in the waters of some “ sulfur springs ” 
and in some volcanic gases. The air near sewers and cess¬ 
pools often contains this gas, since it is one product of the 
decay of organic substances which contain sulfur. The 
albumin in the white part of eggs contains sulfur, and when 
eggs decay, hydrogen sulfide is formed; hence the bad smell. 

213. Preparation of hydrogen sulfide. — The gas is usually 
prepared in the laboratory by the interaction of dilute hydro¬ 
chloric, or sulfuric, acid and ferrous sulfide. The equation 

FeS + 2HC1 = H 2 S + FeCb 

Iron Hydrochloric Hydrogen Iron 

Sulfide Acid Sulfide Chloride 

214. Properties of hydrogen sulfide. — This gas is poison¬ 
ous. A little, even if diluted with air, often produces 


166 


A BRIEF COURSE IN CHEMISTRY 


headache and nausea, and a large quantity of the gas may 
prove fatal. Care should be taken to prevent its escape 
into the laboratory. 

One volume of water dissolves about three volumes of 
hydrogen sulfide gas at ordinary temperatures. The solution 
is often called hydrogen sulfide water, and can be used in¬ 
stead of the gas in many chemical experiments; it has a weak 
acid reaction, and is sometimes called hydrosulfuric acid. 
In terms of the ionic theory the solution contains few ions 
(H+ and S—). It forms salts called sulfides (§ 216). 

Hydrogen sulfide burns with a pale bluish flame, forming 
sulfur dioxide and water. If the supply of air is insufficient, 
combustion is incomplete, and sulfur and water are formed. 
Hydrogen sulfide reduces nitric acid and sulfuric acid; the 
equation for the latter reaction is: — 

H2S + H2SO4 = SO2 + S + 2 H2O 

Hydrogen Sulphuric Sulfur Sulfur Water 

Sulfide Acid Dioxide 

215. Sulfides. — These may be regarded as salts of the 
weak acid hydrogen sulfide, though they are not always 
prepared directly from hydrogen sulfide. They can be 
produced by the direct union of melted sulfur and metals, 
as in the case of iron sulfide previously mentioned (§8), or 
by exposing metals to the moist gas. 

Sulfides are usually prepared in the laboratory by pre¬ 
cipitation. That is, the gas is passed into solutions con¬ 
taining a metallic ion, 6.gr., Cu++; or hydrogen sulfide water 
is added. 

Many sulfides are black, one (zinc sulfide) is white, and 
several have a characteristic color. Thus, arsenic sulfide 
is pale yellow, manganese sulfide is flesh colored, and anti¬ 
mony sulfide is orange red. The color often affords a ready 
means of identifying a sulfide. 

Copper, tin, lead, and silver react readily with hydrogen 
sulfide and are rapidly tarnished by exposure to the gas. 
Silverware turns brown or black, especially in houses heated 
by coal and lighted by coal gas, probably owing to the small 
quantity of hydrogen sulfide from these sources. A brown 


SULFUR AND SULFIDES 


167 


film (silver sulfide) also coats silver spoons which are put 
into mustard, eggs, and some vegetables, such as cauliflower. 
Lead compounds are blackened by this gas, owing to the 
formation of lead sulfide. For this reason buildings painted 
with “ white lead ” paint often become dark, and, similarly, 
oil paintings are discolored. The blackening of paper mois¬ 
tened with a solution of lead nitrate or acetate is the cus¬ 
tomary test for hydrogen sulfide. 

216. Carbon disulfide. — This substance, when pure, is a clear, 
colorless, heavy liquid, with an agreeable odor, but the commercial 
substance is yellow and has an exceedingly offensive smell. It is poison¬ 
ous, vaporizes very readily, and is highly combustible. When the 
vapor burns, the equation for the reaction is: — 

CS 2 + 3 O 2 = CO 2 + 2 SO 2 

Carbon Oxygen Carbon Sulfur 
Disulfide Dioxide Dioxide 

This liquid must be used with care. No flames should be near when 
carbon disulflde is being evaporated or used as a solvent. (See § 49.) 

EXERCISES 

1. Summarize (a) the physical properties and (6) the chemical 
properties of sulfur. 

2. How is sulfur obtained from deposits in the United States? 

3. What is (a) roll sulfur, (6) brimstone, (c) rhombic sulfur, 
(d) monocHnic sulfur, (e) amorphous sulfur? 

4. How is hydrogen sulfide prepared ? State the equation. 

6. Summarize the properties of hydrogen sulfide. 

6 . State one or more tests for hydrogen sulfide. 

7. Write these equations in the ordinary and the ionic form : (a) Lead 
nitrate and hydrogen sulfide form lead sulfide and nitric acid, (b) Cop¬ 
per sulfate and hydrogen sulfide form copper sulfide and sulfuric acid. 

8. Complete and balance (a) H 2 S H-= CuS H-; 

(6) CdCb +- = CdS +-. 

PROBLEMS 

1. Calculate the weight of sulfur in 500 gm. of pure iron disulfide. 

2. Calculate the weight of sulfur in 77 gm. of sulfuric acid containing 
98 per cent of H 2 SO 4 . 

3. Calculate the simplest formula corresponding to (a) S = 39, 
As = 61; (b) S = 29.9, As = 70.1. 


168 


A BRIEF COURSE IN CHEMISTRY 


SUGGESTIONS FOR LABORATORY WORK 

(References are to Newell’s Laboratory Exercises in Chemistry) 

Exercise *42 — Different Forms of Sulfur. 

Exercise *47 — Hydrogen Sulfide (Short Method). 

Exercise S37 — Hydrogen Sulfide — T. 

Exercise S38 — Sulfides — T. 


CHAPTER XVIII 


SULFUR DIOXIDE —SULFUROUS ACID —SUL¬ 
FURIC ACID 

217. Preparation of sulfur dioxide. — Sulfur dioxide is 
formed by burning sulfur in air (or oxygen), thus: — 

S + O 2 = SO 2 

Sulfur Oxygen Sulfur Dioxide 

The gas is also formed when sulfur compounds are burned 
and sulfide ores are roasted, especially such ores as iron 
pyrites (iron disulfide, FeS 2 ) and galena (lead sulfide, PbS). 
It is usually one of the products of the decomposition of 
certain sulfur compounds, e.g., sulfuric acid (H 2 SO 4 ) and 
sulfurous acid (H 2 SO 3 ). 

On an industrial scale the gas is prepared by burning 
sulfur or by roasting metallic sulfides, especially iron disul¬ 
fide (iron pyrites, FeS 2 ) thus: — 

4 FeS2 + 11 O 2 = 8 SO 2 + 2 Fe203 

Iron Disulfide Oxygen Sulfur Dioxide Iron Oxide 

Both of these reactions (sulfur and sulfide) are utilized on 
a large scale in the manufacture of sulfuric acid (§§ 222- 
224). 

In the laboratory two methods of preparation are used. 
(1) If copper and concentrated sulfuric acid are heated 
together, one of the products of the reaction is sulfur dioxide. 
The equation may be written: — 

Cu + 2 H 2 SO 4 = SO 2 + CUSO 4 + 2 H 2 O 

Copper Sulfuric Acid Sulfur Dioxide Copper Sulfate Water 

169 


170 


A BRIEF COURSE IN CHEMISTRY 


(2) More commonly dilute sulfuric (or hydrochloric) acid 
is added slowly to a sulfite. The equation is: — 

Na2S03 + H2SO4 = SO2 + Na2S04 + H2O 

Sodium Sulfuric Sulfur Sodium Water 

Sulfite Acid Dioxide Sulfate 

The sulfite method is safer and more convenient, especially 
for a steady current of the gas. 

Sometimes the gas is obtained from a cylinder of liquid 
sulfur dioxide (Fig. 85). 

218. Physical properties of sulfur dioxide. — Sulfur 
dioxide is a colorless gas. Its odor is suffocating, being the 
well-known odor of burning sulfur. The gas is a 
little more than twice as heavy as air. A liter 
at 0° C. and 760 mm. weighs 2.9 gm. 

Sulfur dioxide is readily liquefied. The liquid 
is a common article of commerce (Fig. 85). (See 
§ 217, end.)^ 

The gas is very soluble in water; at 20° C. 
about forty volumes of gas dissolve in one volume 
of water. This solution is sour and reddens blue 
litmus; it contains sulfurous acid (H2SO3) besides 
sulfur dioxide. 

Sulfur dioxide does not burn and will not sup¬ 
port ordinary combustion. 

Large quantities of sulfur dioxide are used in 
manufacturing sulfites and sulfuric acid (§§ 220 , 
222-224). 

219. Sulfurous acid. — This compound is pre¬ 
pared by bubbling sulfur dioxide through water. Some of 
the gas combines with the water, thus: — 

SO2 + H2O = H2SO3 

Sulfur Dioxide Water Sulfurous Acid 

Sulfur dioxide is often called sulfurous anhydride, because 
it is the anhydride of sulfurous acid. Anhydrides of non- 
metals {e.g., carbon and sulfur) are acid anhydrides, Le., 
non-metallic oxides which unite with water to form acids. 

Sulfurous acid is unstable and decomposes readily into 
sulfur dioxide and water, especially when the solution is 



SULPHUR 

DIOXIDE 


Fig. 85 .— 
Cylinder 
of liquid 
sulfur di¬ 
oxide. 






SULFUR DIOXIDE — SULFUROUS ACID 


171 


heated; solutions of sulfurous acid smell strongly of sulfur 
dioxide. The formation and decomposition may be repre¬ 
sented as a reversible equation, thus: — 

SO 2 + H 2 O H 2 SO 3 

SuKurous acid is readily oxidized. Solutions of the acid, 
if exposed to air, soon give a test for sulfuric acid, which is 
formed by the combining of the sulfurous acid with oxygen 
from the air. Oxidizing agents, such as potassium perman¬ 
ganate, produce this change quickly. Since the potassium 
permanganate is reduced at the same time, we can also 
describe this chemical change by saying sulfurous acid is a 
reducing agent. 

220. Salts of sulfurous acid. — Sulfurous acid forms two 
classes of salts:— the normal and the acid sulfites. When 
the two atoms of hydrogen in sulfurous acid are replaced by 
a metal, the product is a normal sulfite, e.g., Na 2 S 03 , normal 
sodium sulfite. But if only one atom is replaced, the product 
is an acid sulfite, e.g., NaHS 03 , acid sodium sulfite. 

Acids like sulfurous acid, which have two replaceable 
hydrogen atoms, are called dibasic acids, because they form 
salts by interaction with two different portions of a base. 
For example, normal sodium sulfite (Na 2 S 03 ) is formed when 
sulfurous acid is neutralized by sodium hydroxide according 
to this reaction: — 

H 2 SO 3 + 2NaOH = Na 2 S 03 + 2 H 2 O 

Sulfurous Acid Sodium Hydroxide Sodium Sulfite Water 

Whereas acid sodium sulfite (NaHSOs), often called bisulfite 
of soda, is formed when half as much base is used, thus: — 

H 2 S 03 + Na0H = NaHS03 + H 2 O 

Acid Sodium Sulfite 

Both kinds of salts yield sulfur dioxide by interaction with 
an acid (§ 217 (2)). The acid salt is sometimes used as the 
antichlor to remove the excess of chlorine from bleached 
cotton cloth (§ 126). It is also used in tanning, and in the 
manufacture of starch, sugar, and paper. 

The corresponding calcium salt, acid calcium sulfite 


172 


A BRIEF COURSE IN CHEMISTRY 


(Ca(HS 03 ) 2 ), is extensively used in one process of manufac¬ 
turing paper from wood. The chips of wood are “ cooked 
in large vessels with a solution of acid calcium sulfite, which 
dissolves the lignin and leaves the cellulose in the form of 
pulp, which is then made into'paper (§ 398). 

221. Uses of sulfur dioxide and sulfurous acid. — Large 
quantities of sulfur dioxide are used in manufacturing sul¬ 
furic acid (§ 222) and acid calcium sulfite. 

Moist sulfur dioxide, which is really a solution of sul¬ 
furous acid, is used as a disinfectant and a bleaching agent. 
It destroys insects and organisms (or prevents their growth). 
It decomposes organic coloring matter, thereby forming 
colorless compounds. Infected clothing 
and rooms are sometimes fumigated with 
sulfur dioxide; the gas for household use 
may be obtained by burning a “ sulfur 
candle ” (Fig. 86). Dried fruits (e.g.y 
peaches, apricots, and apples), canned 
corn, cherries, and nuts are bleached by 
sulfur dioxide. The solution is used to 
bleach silk, hair, straw, paper, wool, and 
other substances which would be injured 
by chlorine (§§ 126,126). In some cases the bleached article, 
e.g., a straw hat, partially regains its color or becomes yellow. 

The specific effect of the moist gas can be shown by putting 
a wet colored flower into a bottle in which sulfur is burning; 
the flower soon loses its color. The general effect of sulfur 
dioxide on organic matter is seen in localities where consider¬ 
able gas escapes, e.g., near smelters and chemical works. 
Here trees, shrubs, and other kinds of vegetation are blighted 
or destroyed. 

Sulfur dioxide is used as the cooling substance in one kind 
of “ iceless ’’ refrigerators. The liquid flows through coils, 
evaporates, and cools the coils, then passes through a com¬ 
pressor (driven by a small electric motor), becomes liquid 
again, and so on. By evaporation heat is absorbed and 
keeps the air cool in the refrigerator. (Compare § 194.) 

222. Manufacture of sulfuric acid. — Enormous quan¬ 
tities of sulfuric acid are manufactured by two processes, 




Fig. 86. — Sulfur can¬ 
dle for fumigation 
of a room. 






SULFUR DIOXIDE — SULFUROUS ACID 


173 


known as (1) the lead-chamber process and (2) the contact 
process. In each process sulfur dioxide is oxidized to sulfur 
trioxide, which with water forms sulfuric acid. A general 
equation for the essential chemical change is: — 

2 SO 2 + O 2 + 2 H 2 O = 2 H 2 SO 4 

In the lead-chamber process the oxidation is accomplished 
by nitrogen oxides; and in the contact process it is hastened 
by a catalyst. The lead-chamber process is largely used 
when acid of a moderate concentration is needed, e.g., for 
manufacturing phosphate fertilizer (§§ 301, 302). The 
contact process is used to produce concentrated acid. 

223. Manufacture of sulfuric acid by the lead-chamber 
process. — Sulfur dioxide, air, steam, and nitrogen oxides 
are introduced into large lead chambers. These gases 
react and produce sulfuric acid, which collects on the floors 
of the lead chambers. This acid contains 60 to 70 per cent 
of the compound H 2 SO 4 . For some uses, the acid needs 
no further treatment and can be used at once. The acid 
can be concentrated by heating it first in lead-lined pans 
and finally in quartz vessels until it contains about 94 per cent 
of H 2 SO 4 (and has a specific gravity of 1.84). 

There are three main chemical changes in the lead-chamber 
process: (1) Sulfur dioxide is oxidized to sulfur trioxide by 
nitrogen dioxide, thus : — 

SO 2 + NO 2 = SO 3 + NO 

Sulfur Nitrogen Sulfur Nitric 

Dioxide Dioxide Trioxide Oxide 

(2) Sulfur trioxide and water form sulfuric acid, thus: — 

SO3 + H2O = H2SO4 

Sulfur Trioxide Water Sulfvuic Acid 

(3) Nitric oxide from (1) unites with oxygen (from the 
admitted air) and forms nitrogen dioxide (which is used over 
again), thus: — 

2 NO + O2 = 2NO2 

Nitric Oxide Oxygen Nitrogen Dioxide 

Hence, nitric oxide acts as a carrier of oxygen, so to speak. 


174 


A BRIEF COURSE IN CHEMISTRY 


from the (admitted) air to the sulfur dioxide. Some authori¬ 
ties call the nitric oxide a catalyst (§§ 22, 78, 193, 224). 

224. Manufacture of sulfuric acid by the contact pro¬ 
cess. — In this process sulfur dioxide and air, well purified 
and heated to about 400° C., are brought in contact with a 
catalyst, usually platinum. The sulfur dioxide is quickly 
oxidized to sulfur trioxide, thus: — 

2 SO2 “h O2 = 2 SO3 

Sulfur Dioxide Oxygen Sulfur Trioxide 

The sulfur trioxide is conducted into sulfuric acid containing 
a little water (because it is not absorbed quickly enough by 
water alone) and thereby produces sulfuric acid, thus: — 

SO3 + H2O = H2SO4 

Sulfur Trioxide Water Sulfuric Acid 

The essential factor in the contact process is the catalyst. 
It must present a large surface for contact with the gases, be 



Fig, 87. — Sketch of the apparatus for making sulfuric acid by the contact 
process. 


not easily “ poisoned ” by impurities in the gas, and hasten 
the reaction to a velocity which permits profitable operation. 
The best catalyst is platinum, which is coated in a very finely 
divided state on the surface of asbestos fibers, small lumps 
of magnesium sulfate, or silica gel {i.e., silicic acid containing 
5 to 7 per cent of water). Other catalysts may be used 
instead of platinum, e.g., iron oxide or vanadium silicate. 

The construction and operation of a contact acid plant is clear from 
the sketch shown in Fig. 87. The blower A forces air into the 
burner B, where the sulfur dioxide is formed by burning iron pyrites 















































































SULFUR DIOXIDE — SULFUROUS ACID 


175 


(FeS 2 ) or sulfur. The gases pass into the dust chamber C, where they 
are freed from sulfur dust and other solid impurities; this is an impor¬ 
tant step, for dust reduces the transforming power of the catalyst. The 
gases, cooled by the pipe D, are further cleaned in the scrubbers, which 
contain coke wet with water (E) and with sulfuric acid (F). The 
next step is the removal of dust and traces of arsenic compounds (e.g., 
arsenious oxide AS 2 O 3 ) in the purifier G, which would “ poison ” the 
platinum used as a catalytic agent and stop the formation of sulfur 
trioxide. 

The purified gases (mainly sulfur dioxide) then enter the mixer and 
heater H. Here a large excess of air is introduced from the blower 
and the whole mixture is heated to 400°C. This temperature is care¬ 
fully regulated because at 400° C. the yield of sulfur trioxide is maximum 
(98-99 per cent). 

The purified and heated mixture of sulfur dioxide and air passes 
into the contact chamber 7. Here the gases come in contact with the 
catalyst and form sulfur trioxide. The catalyst, if platinum, usually con¬ 
sists of asbestos fibers coated with a very thin layer of metallic platinum 
and is spread out on plates or mixed with porous material in order to 
provide a large contact surface. 

The final step is the transformation of the sulfur trioxide into sulfuric 
acid by combining with water; the trioxide is passed into the absorber 
(not shown). This is a large earthenware jug partly 
filled with sulfuric acid containing 1 to 3 per cent of 
water. In this liquid all the sulfur trioxide combines 
with water; the water is replenished to maintain the 
required concentration in the absorber. The gases 
from the contact chamber I cannot be passed di¬ 
rectly into pure water owing to the formation of a 
mist or fog which prevents absorption of the sulfur 
trioxide. 

If fuming sulfuric acid (H2S2O7), or oleum as it is 
sometimes called, is required, the sulfur trioxide is 
passed directly into 100 per cent sulfuric acid. 

225. Physical properties of sulfuric acid.— 

Sulfuric acid is an oily liquid, colorless when 
pure, though sometimes brown from the pres¬ 
ence of charred organic matter, such as dust. -piG. 88.—Find- 
The specific gravity of the commercial acid is ing the specific 

about 1.84 ; thus it is nearly twice as heavy as fur^acid^^h 

water (Fig. 88). When heated, it begins to a hydrometer, 

decompose and form white, suffocating fumes 
of sulfur trioxide at 150°-180° C., and finally boils at about 
338° C. 







176 


A BRIEF COURSE IN CHEMISTRY 


226. Chemical properties of sulfuric acid. — Sulfuric 
acid mixes with water in all proportions, and during the 
mixing much heat is evolved. The acid should always be 
poured into the water and the mixture should be stirred, 
otherwise the intense heat may crack the vessel or spatter 


the hot acid. 

This tendency to absorb water is shown in many ways. 
The concentrated acid absorbs moisture from the air and 
from gases passed through it. It is often used in the labo¬ 
ratory to dry gases. Organic substances, such as wood, 
paper (cellulose (CeHioOs)^), sugar (C6H12O6 and C12H22O11), 
starch ((CeHioOs)®), and cotton, are charred by concentrated 
sulfuric acid (Fig. 89 ). Such compounds contain hydrogen 
and oxygen in the proportion of 2 H to 10 , 
i.e., in the proportion to form water; hence 
these two elements are abstracted and carbon 
alone remains. Sulfuric acid also disinte¬ 
grates the flesh, often causing serious bums. 
If accidentally spilled on the hands or spat¬ 
tered on the face, it should be washed off 
immediately. 

The interaction of sulfuric acid and metals 
varies. With many metals dilute sulfuric 
acid forms hydrogen and the corresponding 
metallic sulfate (§ 68 ). Thus, hydrogen is 
usually prepared in the laboratory from zinc and sulfuric 
acid. Some metals, however, reduce the acid to sulfur dioxide 
(or hydrogen sulfide). Thus, the equation expressing the 
reaction with copper and concentrated acid is: — 



Fig. 89. — Paper 
charred by 
concentrated 
sulfuric acid. 


Cu + 2 H2SO4 = CUSO4 + SO2 + 2 H2O 

Copper §uIfTiric Copper Sulfur Water 

Acid Sulfate Dioxide 


Iron is the only common metal that is not readily attacked 
by the concentrated acid, and advantage is taken of this 
property in transporting the acid in bulk in iron tank 
cars. 

Sulfuric acid unites with ammonia (NH3) to form am¬ 
monium sulfate ((NH4)2S04). 




SULFUR DIOXIDE — SULFUROUS ACID 177 

Hot concentrated sulfuric acid and carbon form sulfur 
dioxide and carbon dioxide. Thus: — 

2 H 2 SO 4 + C = 2 SO 2 + CO 2 + 2 H 2 O 

Dilute solutions of sulfuric acid contain an abundance of 
hydrogen ions (H+) and sulfate ions (SO4 —). The HSO4- 
ion is in solutions of certain concentration. 

227. Uses of sulfuric acid. — Sulfuric acid is one of the 
most important substances. Directly or indirectly it is 
used in hundreds of industries upon which the comfort, 
prosperity, and progress of mankind depend. It is used 
in the manufacture of many acids (§§ 130, 198). The 
petroleum industry requires large amounts for refining the 
oil — about 20 per cent of the output. Enormous quantities 
are consumed in making fertilizers, alum and other sulfates, 
nitro-glycerin, glucose, dyes, and in various parts of such 
fundamental industries as dyeing, bleaching, metal cleaning 
{e.g., pickling iron castings), refining, and metallurgy. 

228. Sulfates. — Sulfuric acid is dibasic and forms two 
classes of salts — the normal sulfates, such as sodium sulfate 
(Na 2 S 04 ), and the acid sulfates (or bisulfates), such as acid 
sodium sulfate (NaHS 04 ). (Compare § 220.) 

Most sulfates are soluble in water; only the sulfates of 
barium, strontium, and lead are insoluble, while calcium 
sulfate is slightly soluble. Important sulfates are calcium 
sulfate (gypsum, CaS 04.2 H 2 O), barium sulfate (barite, 
barytes, heavy spar, BaS 04 ), zinc sulfate (ZnS 04 , and white 
vitriol, ZnS 04 .7 H 2 O), copper sulfate (CUSO 4 , and blue vit¬ 
riol or blue stone, CUSO 4 .5 H 2 O), iron sulfate (ferrous sulfate, 
FeS 04 , and green vitriol, copperas, FeS 04 .7 H 2 O), sodium 
sulfate (Na 2 S 04 ) and Glauber^s salt (Na 2 S 04.10 H 2 O), and 
magnesium sulfate (MgS 04 , and Epsom salts, MgS04. 7 H 2 O). 

229. The test for sulfuric acid or for a soluble sulfate. — 
The usual test is the formation of white, insoluble barium 
sulfate upon the addition of barium chloride (§ 178). 

An insoluble sulfate, e.g., calcium sulfate or barium sulfate, 
if fused on charcoal is reduced to a sulfide, which blackens 
a moist silver coin, owing to the formation of silver sulfide 
(Ag 2 S); this is the usual test for an insoluble sulfate. 


178 


A BRIEF COURSE IN CHEMISTRY 


230. Sodium thiosulfate (Na 2 S 203 ). — This is a salt of 
an unstable acid. It is sometimes incorrectly called sodium 
hyposulfite, or simply ‘‘ hypo.’’ It is a white, crystalline 
solid, very soluble in water. The solution, used in excess, 
dissolves certain compounds of silver, i.e., AgCl, AgBr, Agl; 
hence its extensive use in photography (§ 491). 

EXERCISES 

1. How is sulfur dioxide prepared? Give four equations for its prep¬ 
aration. 

2. State the properties of sulfur dioxide. 

3. Define, and give the formula of, (a) a normal sulfite and (&) an 
acid sulfite. 

4. Apply Exercise 3 to sulfates. 

6. Describe the contact process of manufacturing sulfuric acid. 

- 6. Enumerate the important uses of suKuric acid. 

7. What is (a) gypsum, (b) white vitriol, (c) green vitriol, (d) blue 
vitriol, (e) Glauber’s salt, (f) oil of vitriol, (g) “ hypo,” (h) calcium 
bisulfite ? 

8. State the test for (a) sulfuric acid, (b) sulfurous acid, (c) a sol¬ 
uble sulfate, (d) an insoluble sulfate, (e) a sulfite. 

9. Write the formula of (a) the sulfite, (b) the acid sulfite, (c) the 
sulfate, (d) the acid sulfate of NH4, Ca, Pb, Ag, Ba, Zn. 

PROBLEMS 

1. What weight of sulfur dioxide can be prepared from 25 gm. of 
sodium sulfite (92 per cent pure)? 

2. How much sulfur (99 per cent pure) is needed to manufacture 100 
tons of sulfuric acid containing 5 per cent of water? 

3. A flask filled with water was found to weigh 72 gm., the flask alone 
weighing 22 gm. The flask filled with sulfuric acid weighed 114 gm. 
Calculate the specific gravity of the sulfuric acid. 

4. What weight of pure sulfuric acid can be made from (a) 1000 tons 
of pure sulfur and (b) 1000 tons of pure FeS 2 ? 

SUGGESTIONS FOR LABORATORY WORK 
(References are to Newell’s Laboratory Exercises in Chemistry) 

Exercise *43 — Sulfur Dioxide (Short Method). 

Exercise 44 — Sulfur Dioxide and Sulfurous Acid — T. 

Exercise *45 — Properties of Sulfuric Acid. 

Exercise *46 — Test for Sulfuric Acid, Sulfates, SO^ions. 


CHAPTER XIX 


FUELS—ILLUMINANTS —PETROLEUM AND ITS 
PRODUCTS — DISTILLATION OF COAL —FLAMES 

231. Carbon and energy. — When carbon and many 
of its compounds burn, the chemical energy is transformed in 
part into heat energy or light energy. Hence carbon and 
some of its compounds are fuels and illuminants. 

232. Fuels and illuminants. — Fuels are substances which 
are burned to produce heat, e.g., (a) wood, charcoal, coke,' 
and coal, (6) petroleum, fuel oil, kerosene, and gasolene, and 
(c) natural, producer, water, and coal gases. 

Illuminants are substances which are burned to produce 
light, e.g., coal gas, oil gas, acetylene (gas), and kerosene. 

233. Composition of fuels. — Wood is composed mainly 
of cellulose (a compound of carbon, hydrogen, and oxygen 
(§ 396)), water (10 to 50 per cent), and a small per cent of 
mineral matter. Charcoal and coke are nearly pure carbon 
(43, 44). Hard coal is about 90 per cent carbon, and soft 
coal about 70 per cent. All kinds of coal contain mineral 
matter, which is left as ashes when the coal is burned. Soft 
coal contains moisture and volatile matter. 

Liquid fuels (except alcohol) are mixtures of hydrocarbons 
(compounds of hydrogen and carbon, § 246). Alcohol is a 
compound of carbon, hydrogen, and oxygen. The two 
common kinds of alcohol are ethyl alcohol or ethanol 
(C2H5OH) and methanol (CH3OH). (See §§ 265, 266.) 

Gaseous fuels contain hydrocarbons, hydrogen, and carbon 
monoxide; acetylene is a hydrocarbon (C 2 H 2 ). 

234. Combustion of fuels. — Fuels bum when heated to 
the proper temperature. Chemically this means that the 
carbon, hydrogen, and carbon monoxide (in the fuels) unite 

179 


180 


A BRIEF COURSE IN CHEMISTRY 


with oxygen furnished by air. The gaseous products of com¬ 
bustion are carbon dioxide, carbon monoxide, and water. 
The solid products are the ashes and, if combustion is incom¬ 
plete, smoke. 

235. Measurement of heat produced by fuels. — The es¬ 
sential characteristic of a good fuel is its heat-producing 
capacity, especially when used to produce steam. Heat, 
like other forms of energy, is measured by a special unit. 
The unit used for fuels is the British thermal unit (B. t. u.). 
It is the amount of heat that will raise the temperature of 
1 pound of water 1 degree measured on the Fahrenheit ther¬ 
mometer, or briefly 1 B. t. u. = 1 lb. water raised 1° F. The 
number of heat units produced by a given quantity of fuel, 
usually 1 pound or 1 cubic foot, is called its fuel value. 
Thus, coke yields about 14,300 B. t. u. per pound. Fuel 
oil, now extensively used in place of coal, yields about 18,500 
B. t. u. per pound. Fuel gases (producer, water, and coal 
gas) give 145 to 600 B. t. u. per cubic foot. 

Besides the British thermal unit, there is another heat 
unit called the small calorie (cal.). It is the amount of heat 
required to raise the temperature of 1 gram of water 1° C. 
(usually 15° to 16°). One B. t. u. equals 252 cal. 

236. Thermo-chemical equations. — The heat liberated 
when carbon, or a carbon compound, burns can be incorpo¬ 
rated in the equation expressing the chemical change. 
Thus, the thermo-chemical equation for the burning of car¬ 
bon in the form of pure charcoal is: — 

C + O 2 = CO 2 + 97,000 cal. 

In this equation C stands for 12 gm. of carbon, O 2 for 32 
of oxygen, CO 2 for 44 of carbon dioxide, and cal. for small 
calories. 

237. Burning of coal. — Coal is the commonest fuel. 
Its fuel value varies with the kind and quality, e.g., bitu¬ 
minous or soft coal yields from 9000 to 14,700 B. t. u. (or 
8700 cal.) per pound and anthracite from 12,500 to 13,700. 

When coal is burned, the products of combustion depend 
on the kind of coal and on the proportion of air supplied. 
With an excess of air, hard coal and good soft coal produce 


FUELS — ILLUMINANTS — PETROLEUM 181 


carbon dioxide and water. If too much air is supplied, heat 
is also lost up the chimney.’^ If the air is insufficient, car¬ 
bon monoxide is formed, and some unburned carbon escapes 
as smoke, especially with soft coal. 

The air needed for burning coal in stoves or under boilers 
is drawn in through an opening under the grate by the draft 
created by the gases as they rise up the chimney, though 
sometimes the air is forced in by a ‘‘ blower. 

Both carbon dioxide and carbon monoxide are formed in a 
coal fire in a stove. When air enters the bottom of the grate 
(Fig. 90) and comes in contact with the burning coal, oxygen 
unites with carbon to form carbon dioxide. The carbon 
dioxide rises through the hot coal, and the carbon reduces the 




CO 2 + C = 2 CO 


C + O 2 = CO 2 


Fig. 90. — The three main changes during combustion in a coal fire. 

carbon dioxide to carbon monoxide. Some of the carbon 
monoxide usually escapes, but most of it burns with a flicker¬ 
ing bluish flame on the top of the fire, forming carbon dioxide 
with the oxygen of the air. 

The combustion of coal in a stove or furnace in a house is 
regulated by dampers, usually three — a lower one in the 
door below the grate (A), a middle one in the door just above 
the fire (B), and an upper one in the pipe connected with the 
chimney (C) (Fig. 90). 

When the fire is built, or needs to be started up,” the 
dampers in the lower door and the pipe are opened but the 
one in the upper door is closed. This arrangement allows 
plenty of air to pass up through the burning coal and increase 
the combustion, and also creates a draft by permitting the 
hot products of combustion to rise and escape out of the chim- 



















182 


A BRIEF COURSE IN CHEMISTRY 


ney. Once started, the combustion in the fire can be regu¬ 
lated by closing the lower damper and opening the middle 
one, partly or wholly, and adjusting the chimney damper; 
by this arrangement most of the air goes over the fire instead 
of up through it. 

Care must be taken to admit enough air through the 
lower door to burn the coal as completely as possible, so that 
the maximum quantity of heat will be liberated. Special 
care should also be taken to prevent the escape of “ coal gas 
into the house. This gas contains carbon monoxide, which 
is poisonous (§ 62). 

238. Burning other solid carbon fuels. — Charcoal burns 
with a slight fiame, and yields no smoke. Its fuel value is 
7000 to 8000 B. t. u. per pound. Coke also bums with a 
small fiame and without smoke. Its fuel value is about 
14,300 B. t. u. per pound. Both charcoal and coke find 
extensive use in the iron and steel industry (§§ 305, 312). 

Wood has been used as a fuel for ages. The hard varieties 
such as oak, ash, and maple are the best fuels. Dry wood 
yields from 5600 to 8000 B. t. u. per pound. 

239. Alcohol fuels. — Different kinds of alcohol are used 
as fuel, usually on a small scale. They burn without smoke 
and have a high fuel value. Methyl alcohol (CH3OH), also 
called wood alcohol (§ 265) and methanol, gives about 9600 
B. t. u. per pound. It is the fuel in “ solid alcohol ” (§ 265). 
Ethyl alcohol (C2H5OH), also called grain alcohol or simply 
alcohol, gives about 12,700 B. t. u. per pound. 

240. Fuels from petroleum. — Three products from petro¬ 
leum — fuel oil, gasolene, and kerosene — are used as fuels. 
They are mixtures of hydrocarbons (§ 246). In burning, 
these compounds are decomposed; the carbon forms carbon 
dioxide and the hydrogen forms water. They produce hot 
flames. 

Fuel oil must be supplied to the furnace in the form of 
spray. This is produced by forcing the oil through a fine 
opening or by blowing it with steam or air through an 
atomizer. When the spray of oil burns, a large amount of 
heat is liberated, as high as 19,800 B. t. u. per pound in the 
case of some grades. 


FUELS — ILLUMINANTS — PETROLEUM 183 


The general principle on which the latter type of burner operates 
is shown in Fig. 91. Oil enters at A and flows through D into the mixing 
and atomizing chamber C. Steam enters at B and passes through F 
and F into the chamber C, where it cuts across the oil at an angle. 
Here the oil and steam mix, and the pressure forces the oil as a fine spray 
out into the furnace, where it burns instantly with a hot flame. 



Fig. 91. — Fuel oil burner. 


Fuel oil is extensively used on warships and steamships, 
and in many manufactories. By using oil in place of coal to 
generate steam, additional space is provided for storage and 
more efficient combustion is attained. 

Gasolene is the fuel used in the engines of automobiles, 
trucks, motor boats, motor cycles, airships, and airplanes. 
Gasolene is very volatile and 
the vapor burns readily. If 
the vapor is mixed with air 
and the mixture is ignited by 
an electric spark, the com¬ 
bustion is so rapid that it is 
practically an explosion; the 
gases, suddenly expanded, ex¬ 
ert pressure, which is converted by the machinery into steady 
and continuous motion. 

Kerosene is used as a fuel to a limited extent in engines 
and in cooking stoves. In some portable stoves kerosene 
is burned by means of a large wick (as in a lamp). The 
fuel value is sometimes as high as 19,900 B. t. u. per pound. 

241. Petroleum and its products. — Petroleum is an oil 
which is distilled to obtain gasolene, kerosene, lubricating 
oil, vaseline, and paraffin. Its composition is complex, but 
all varieties are essentially mixtures of liquid and solid hy¬ 
drocarbons (§ 246). Certain grades contain compounds of 
nitrogen and of sulfur. 

In some localities the oil issues from the earth, but it is 
usually necessary to erect a derrick, drill a deep hole, and 
insert a pipe into the porous rock containing the oil. Some¬ 
times the oil “ spouts ” out of the well when first drilled, 
but after a time a pump is usually needed to draw it to the 
surface. 

The oil is forced by powerful pumps through large pipes 












184 


A BRIEF COURSE IN CHEMISTRY 


to central points for storage, or more often many miles to the 
refinery, where the petroleum is separated into various com¬ 
mercial products by refining. 

242. Refining petroleum. — The petroleum is first cleaned 
by settling and filtration. Then it is distilled by heating it in 
huge retorts (Fig. 92), condensing the vapors in coiled pipes 
immersed in cold water, and collecting the distillates in sepa¬ 
rate tanks. This process is called fractional distillation. 
Certain products, e.g., fuel oil, are obtained from the residue 
left in the stills (§ 244). 

The different distillates are further separated and purified 
by redistillation (Fig. 92, top). The final products depend on 
the composition of the original petroleum and also on the 
demand for special commercial mixtures, e.g., gasolene. 
The chief commercial liquid products are gasolene (boiling 
point 60°-190° C.), kerosene (150°-250° C.), fuel oil (250°- 
350° C.) and various grades of lubricating oils (above 300° C.). 

Formerly the gasolene fraction from ordinary distillation 
of petroleum met the limited need. The first step taken to 
increase the supply was cracking, i.e., the higher boiling oils 
were vaporized and the vapor heated to a high temperature 
(350°-450° C.) and under increased pressure (4 to 5 atmos¬ 
pheres). By this treatment, complex hydrocarbons decom¬ 
posed and formed hydrocarbons within the gasolene range 
(§ 246). Another step was blending, i.e., mixing the higher 
boiling hydrocarbons formerly sold as kerosene with “ casing¬ 
head gasolene.^' The latter contains chiefly the volatile 
hydrocarbons, pentane and hexane, and is obtained from 
natural gas by several methods, one of which is cooling and 
compressing the gas (to 15 to 20 atmospheres). The gaso¬ 
lene sold at present is usually blended and has a wide range 
of boiling point. 

243. Kerosene. — This liquid is a mixture of hydrocar¬ 
bons which have a higher boiling point than those in gasolene. 
Owing to its extensive use as an illuminant, crude kerosene 
is carefully freed from readily combustible liquids and gases, 
which might cause an explosion, and also from tarry matter 
and semi-solid hydrocarbons which would clog a wick. The 
purification is done by agitating the crude kerosene succes- 


FUELS — ILLUMINANTS — PETROLEUM 185 




Fiq. 92 . — Refining petroleum. Fire stills (bottom), steam stills (top), 
agitators (middle). 


















186 


A BRIEF COURSE IN CHEMISTRY 


sively with sulfuric acid, sodium hydroxide, and water (Fig. 
92, middle). 

Commercial kerosene must have a legal flashing point. 
This is the temperature at which the oil gives off sufficient 
vapor to form a momentary flash when a small flame is 
brought near its surface.’^ The legal minimum flashing point 
in most states is about 110° Fahrenheit (about 44° C.). 

244. Fuel oil and other products. — The oil left in the fire 
still (Fig. 92, bottom) after the removal of the low-boiling 
liquids including gasolene, and sometimes kerosene, is fuel oil. 



Fig. 93. — Sketch of a plant for the separation of gasolene and fuel oil from 
petroleum. 


Other oil residues, and crude petroleum itself, are also used 
as fuels under the name fuel oil. Specifications usually pre¬ 
scribe the properties of the fuel oil required, especially its 
heating value, freedom from water and from sulfur com¬ 
pounds, and mobility. 

From the residue after the distillation of the kerosene 
many grades of lubricating oil are obtained; also vaseline 
and paraffin wax. Petroleum oils yielding these products 
are said to have a paraffin base; Pennsylvania and Ohio 
oils are examples. Vaseline finds extensive use as an oint¬ 
ment. Paraffin wax is made into candles and into a water¬ 
proof coating for many substances. The final residue in 















































































FUELS — ILLUMINANTS — PETROLEUM 187 


the still is mainly carbon and is called petroleum coke; 
it is made into electric light carbons. 

Some oils, such as those from Mexico and California, leave 
a thick, black pitch. These oils are said to have an asphalt 
base. 

246. Separation of fuel oil and gasolene from petroleum. — The 

steps in the separation of gasolene and fuel oil from petroleum are shown 
in Fig. 93. This figure also illustrates the general process of refining. 
The crude oil is pumped from the storage tank A into the still B (Fig. 92, 
bottom), which is heated on the bottom by a flame. The vapors pass 
into the condenser C, and the distillate flows into the receiving house D, 
where it is examined through look boxes. The different portions (called 
fractions) are run into the proper tank. After certain low-boiling hydro¬ 
carbons (making up gasolene and kerosene chiefly) have been boiled off 
in the fire still, the residue is drawn off and stored in the fuel oil tank E. 
The benzine fraction F is pumped to the agitator G (Fig. 92, middle) 
and then to the storage tank JI. From here the hquid goes to the still 
/, which is heated by steam (Fig. 92, top). The vapors are condensed 
in /, the distillate received in K, examined, and directed into the proper 
tank — the main one being the gasolene tank L. 

246. Hydrocarbons. — These are compounds of hydrogen 
and carbon. They occur as gases, liquids, and solids. Over 
two hundred and fifty are known. Many are found in petro¬ 
leum, natural gas, asphalt, and coal tar. Hydrocarbons 
are divided into series according to their composition. The 
commonest series is the methane or paraflSn series. The 
first members in order are methane (CH4), ethane (C 2 H 6 ), 
propane (CsHg), butane (C4H10), pentane (C5H12), and 
hexane (C 6 H 14 ). 

Petroleum is a mixture of many hydrocarbons, and most 
American varieties consist largely of paraffin hydrocarbons. 
Gasolene is mainly a mixture of hexane (CeHu), heptane 
(CyHie), and octane (CsHis). Kerosene consists of hydro¬ 
carbons which are composed of ten to sixteen carbon atoms, 
vaseline twenty-two and twenty-three, and paraffin wax still 
higher. 

247. Gaseous fuels. — These include natural gas and the 
various mixtures obtained from coal, e.g., producer gas 
(§§ 66, 386), water gas (§§ 66, 386, 387), and coal gas. 

Natural gas exists in the earth in Pennsylvania, Ohio, 


188 


A BRIEF COURSE IN CHEMISTRY 


West Virginia, and other parts of the United States, usually 
in regions where petroleum is found. It contains from 90 to 
98 per cent methane (CH 4 ), which is the chief heat-producing 
constituent. Natural gas burns with a hot flame. The fuel 
value is about 1000 B. t. u. per cubic foot. It is used as a 
fuel to heat houses and to generate steam in many industries, 
e.g., making steel, glass, brick, and pottery. 

248. Coal gas. — This is made by dry or destructive dis¬ 
tillation of coal, i.e.j by heating coal to a high temperature in 
closed retorts. The volatile product (§ 250) is largely a 
mixture of hydrogen (45 to 50 per cent) and methane (30 to 
36 per cent); other ingredients are hydrocarbons (ethylene, 
etc.), carbon monoxide (6 per cent) and dioxide, and nitrogen. 
It burns with a yellow flame. When used as a fuel, consider¬ 
able air is admitted and a 
special burner is used. In 
the laboratory we use the 
Bunsen burner (§ 252). 
In houses the gas range 
burner is merely several 
small modifled Bunsen 

burners arranged advantageously (Fig. 94). The fuel value 
of coal gas varies from 525 to 600 B. t. u. per cubic foot. 

249. Illuminants. — Some of the gases used as fuel are 
also used as a source of light, i.e., as illuminants. Coal gas 
and enriched water gas are the most important (§§ 65, 386, 
387). 

250. Manufacture of coal gas. — Coal is subjected to dry 
distillation to prepare coal, or illuminating, gas. The gas 
is used as a fuel, e.gr., in a kitchen gas range, and also as an 
illuminant. 

The manufacture of coal gas can be understood from the diagram of 
a coal gas plant shown in Fig. 95. (1) The coal is heated externally for 

several hours in closed retorts (A). The volatile products escape from 
the retorts and bubble through water into a large iron pipe called the 
hydraulic main (B). Here some of the tar is deposited and ammonium 
compounds are dissolved by the water which flows constantly through 
the main and prevents the gas from escaping back into the retorts. 
The ammoniacal hquor and tar flow into the tar well (C). 

(2) From the hydrauHc main the hot and impure gas passes through 


Air 



Air 


Fig. 94. — A gas range burner. 








FUELS — ILLUMINANTS — PETROLEUM 189 


vertical pipes called the condenser (D), which cools the gas and removes 
tar. 

(3) An exhauster {E) draws or forces the gas from the condenser 
into the scrubber (and onward through the purifiers into the gas holder). 

(4) The scrubber (F) is a tall tower filled with coke or pebbles over 
which ammoniacal liquor and water trickle. The object of the scrubber 
is to remove the remaining ammonium compounds, and also carbon 
dioxide, hydrogen sulfide, and the last traces of tar. 

(5) From the scrubber the gas passes into the purifier (G), which is a 
series of shallow rectangular boxes fitted with frames loosely covered 
with hme or hydrated ferric oxide, or both. In the purifier the remain¬ 
ing carbon dioxide and sulfur compounds are removed. 

(6) The purified gas next passes through a large meter, which records 
its volume, and then into a gas holder ( H), where it is stored over water 
and finally forced through pipes to the consumer. 



A ton of good gas coal yields about 10,000 cubic feet of 
gas, 1400 pounds of coke (§ 44), 120 pounds of tar, 20 gal¬ 
lons of ammoniacal liquor (§ 189), and a varying amount of 
gas carbon (§ 45). 

The tar, or coal tar as it is often called, collected from the 
hydraulic main and condenser, is a thick, black, foul-smelling 
liquid. Some is used for preserving timber, making road 
material, tarred paper, and black varnishes, and as a protec¬ 
tive paint. Most of it is separated by distillation into its 
important constituents. Among these are the hydrocarbons 
benzene (CeHe), toluene (C 7 H 8 ), naphthalene (CioHg), and 





































































190 


A BRIEF COURSE IN CHEMISTRY 


anthracene (CuHio). They are indispensable substances, 
being the parent substances of dyes, medicines, photographic 
chemicals, and explosives. Naphthalene is sometimes called 
moth balls and is used as a substitute for camphor. Among 
other products from coal tar are creosote oils and 
carbolic acid (phenol, CeHsOH), which are useful 
disinfectants. 

261. Illuminating gas. — Coal gas is often 
burned alone but water gas is usually mixed with 
60 or 70 per cent of coal gas. This mixture is 
called “ illuminating gas.” Owing to the high 
percentage of carbon monoxide, water gas and 
mixtures containing it are poisonous (§ 62). 

Illuminating gases are mixtures of varying pro¬ 
portions. Table VI shows the approximate com¬ 
position. 

Methane, hydrogen, and carbon monoxide burn 
with a feeble (non-yellow) flame; they furnish 
heat, but little light. The illuminants consist of 
ethylene (C2H4), acetylene (C2H2), benzene (CeHe), 
and other hydrocarbons. The illuminants furnish 
the carbon particles which give the 
— flame its yellow color. 

252. The Bunsen burner and its 
flame. — When a mixture of air and 
r.G. 96.-Parte of atypical iUummating gas is burned in a suit- 
Bunsen burner. able burner, a flame is produced 



TABLE VI. — Composition of Illuminating Gases 


Constituents 

Coal Gas 

Water Gas 

Mixed Gas 

Methane. 

34.44 

15.23 

29.81 

Illuminants. 

4.17 

11.07 

6.38 

Hydrogen. 

50.31 

35.04 

46.51 

Carbon Monoxide. 

6.60 

33.42 

12.02 

Carbon Dioxide. 

0.81 

2.66 

1.78 

Nitrogen. 

3.67 

2.24 

3.28 

Oxygen. 

0.00 

0.34 

0.22 


























FUELS — ILLUMINANTS — PETROLEUM 191 


which is non-luminous and hot, and deposits on carbon. 
Such a flame is called the Bunsen flame. It was first pro¬ 
duced in a burner devised by the German chemist Bunsen. 
This burner is used in laboratories as a source of heat. The 
parts of a typical burner are shown in Fig. 96. 

The gas enters the base and flows from a small opening into the long 
tube, which screws down over this opening. At the lower end of the 
tube and above the inlet are two holes, through 
which air is drawn by the gas as it rushes out of the 
small opening. The gas and air mix as they rise 
in the tube, and the mixture burns at the top of the 
long tube. 

The size of the air holes at the bottom can be 
changed by a movable ring. When the holes are 
wide open, the typical non-luminous Bunsen flame 
is formed; this flame is free from soot, and appa¬ 
ratus heated by it is not blackened. When the 
holes are closed, the gas burns with a luminous 
flame and deposits carbon. 

The gas burns at the top of the tube and not 
inside, because the proper mixture of gas and air 
rushes out more quickly than the flame can travel 
back through the tube to the small inlet. If the 
gas supply is slowly decreased, the flame becomes 
smaller, disappears down the tube with a slight ex¬ 
plosion C‘ strikes back ”)> burns at the small gas 
inlet inside the tube. A sudden draft of air, too Fig. 97. — Sketch 
much air admitted through the holes at the lower of a Bunsen 

end of the tube, or too low gas pressure may cause flame, 

the flame to “ strike back.’' 

This modified flame, which has a pale color, a disagreeable odor, and 
deposits soot, should be extinguished and the proper flame produced 
before further use. 

The Bunsen flame (Fig. 97) consists essentially of two 
cones, which may often be distinguished by the different tints. 
The lower and inner cone consists of air and unburned gas. 
It is bluish, but becomes green when too much air is admitted 
(best seen in an imperfect flame). The outer cone is the flame 
proper and consists of burning gases. It is faint blue and 
hot (about 1500° C.). 

The inner cone consists of unburned gases. This can be shown by 
putting one end of a small bore glass tube into the cone; gas will rise 











192 


A BRIEF COURSE IN CHEMISTRY 


through the tube and can be ignited at the upper end (Fig. 98, left). If 
a match is supported by a pin across the top of an unlighted burner, it 
will not become ignited until some time after the gas is first lighted 
(Fig. 98, center). 

The outer cone consists of burning gases. Thus, a match held near 
the flame ignites quickly, while a match laid for an instant across the 
top of the tube is charred only at the two points where it touches the 
outer cone. 

Finally, a wire ga\ize, if pressed down upon the flame, shows a dark 
disk surrounded by a luminous ring due to the 
inner and outer cone respectively 
(Fig. 98, right). 





Fig. 98. — The cones of a Bunsen flame. Drawing unburned gases from 
the inner cone (left). A match does not ignite in the inner cone (center). 
The inner cone produces a dark disk and the outer cone a luminous ring 
(right). 


r 


253. Oxidizing and reducing flames. — The outer portion 
of the Bunsen flame is the oxidizing flame. The inner and 
_ _ lower portion is the re¬ 
ducing flame. A sketch 
of these flames is shown 
in Fig. 97. A is the most 
effective part of the oxi- 
the reducing flame. At A metals 

Fig. 100. — A blow- 


Fig. 99. — A mouth blowpipe. 


dizing flame, and B of 
are oxidized, and at B oxygen compounds 
are reduced. 

Sometimes the oxidizing and reducing 
flames are produced by a mouth blowpipe 
(Fig. 99). A special tube with a flattened 
top is put inside the burner tube to pro¬ 
duce a luminous flame, and the tip of the 
blowpipe is put in or very near this flame, 
and continuously blown through the blowpipe, a long slender 
flame is produced*, called a blowpipe flame (Fig. 100). 


pipe flame. A is 
the oxidizing part 
and B is the reduc¬ 
ing part. 

If air is gently 








FUELS — ILLUMINANTS — PETROLEUM 193 


EXERCISES 

1. What is the fuel value of (a) hard coal, (b) fuel oil, (c) coke ? 

2. Define and illustrate a thermal equation. 

3. Prepare a summary of (a) burning coal, (b) manufacturing coal 
gas, (c) fuel oil. 

4. Explain the use of gasolene in an automobile. 

6. Prepare an outline of the process of refining petroleum. 

6. Complete: (a) C +- = CO; (b) CO +- = CO 2 ; 

(c) CO 2 +- = CO. 

7. Draw a diagram of the apparatus for manufacturing coal gas. 

PROBLEMS 

1. A candle weighing 50 gm. consists of wax composed of 88 per cent 
carbon and 12 per cent hydrogen. What weight of (a) carbon dioxide 
and of (b) water will be formed by burning half the candle? 

2. What volume of air (containing 21 per cent of oxygen by volume) 
will be required for the combustion of 100 tons of coal, assuming that the 
coal is 80 per cent pure carbon and burns to carbon dioxide ? 

3. Ten tons of coke were burned and only 35 tons of carbon dioxide 
were produced. Calculate the per cent of carbon in the coke. 

SUGGESTIONS FOR LABORATORY WORK 

(References are to Newell’s Laboratory Exercises in Chemistry) 

^ Exercise *48 — Distillation of Soft Coal. 

Exercise *49 — Distillation of Wood — T. 

Exercise *50 — Illuminating Gas Flame — T. 

Exercise *51 — Candle Flame — T. 

Exercise *52 — Bunsen Burner Flame — T. 

Exercise *53 — Reduction and Oxidation with the Blowpipe. 

Exercise S39 — Carbonic Acid — T. 

Exercise S40 — Acetylene — T. 

SUPPLEMENTARY SECTIONS FROM PART II 

383. Measurement of fuel value. 

384. The calorimeter is used to find fuel value. 

386. Producer gas. 

386. Water gas. 

387. Manufacture of water gas. 

388. The illuminating gas flame. 

389. Other luminous flames. 

390. The candle flame. 

391. Acetylene and its flame. 

392. The oxy-acetylene flame. 


CHAPTER XX 


SUGAR—STARCH—ACETIC ACID—METHANOL— 
ETHYL ALCOHOL 

264. Sugars. — The most important sugars are (1) ordi¬ 
nary sugar, also called cane sugar and sucrose, (2) dextrose 
or glucose, (3) levulose or fruit sugar, (4) lactose or milk 
sugar, and (5) maltose. 

265. Sucrose. — Sugar cane contains about 18 per cent 
and sugar beets from 12 to 15 per cent; these are the main 
sources of sucrose. 

Sucrose is a white solid. It is very soluble in water; one 
part of water dissolves about three times its weight of sugar 
at ordinary temperatures. If sugar is carefully heated, it 
melts. As the temperature is raised, the sugar begins to 
decompose, water is given off, and a light brown substance 
called caramel is formed which is used to color soups, gravies, 
and beverages. By further heating, a black porous mass of 
carbon is finally obtained, often called sugar charcoal. Thus 
we see that sucrose consists of carbon, hydrogen, and oxygen. 
Its formula is C 12 H 22 O 11 . 

If sucrose is boiled with very dilute acid, it undergoes 
hydrolysis, ^.e., it interacts with water. By this reaction a 
mixture of dextrose and levulose is formed, which is called 

invert sugar. 

Sugar is extracted from sugar cane or sugar beet by treating with 
water and purifying substances, and evaporating the solution until the 
sugar will crystalhze. The crystals are separated by machinery, and 
molasses is left. The raw sugar, as it is called, is refined, i.e., purified 
and crystallized. One important step in the purification is the removal 
of coloring matter by passing the straw-colored solution through a deep 
layer of animal charcoal (§ 43). 


194 


SUGAR — STARCH — ACETIC ACID 


195 


266. Dextrose. — This sugar is about three fifths as sweet 
as sucrose. It is very soluble in water. Dextrose is found 
in honey and in many fruits, especially grapes, and is some¬ 
times called grape sugar. Another name for it is glucose. 

Dextrose ferments with ordinary yeast, thus 

C 6 H 12 O 6 = 2 CO 2 "h 2.C2II60 

Dextrose Carbon Dioxide Alcohol 

Dextrose is an inexpensive substitute for sucrose, and is 
extensively used in making candy, jellies, and sirups. 

257. Levulose. — This sugar is found in fruits and honey, 
and is often associated with dextrose. It is sometimes called 

fructose or fruit sugar. 

258. Testing for dextrose. — Dextrose and levulose are 
reducing agents, and are called reducing sugars. An alkaline 
solution of dextrose is sometimes used to reduce a silver solu¬ 
tion and deposit the silver as a bright film in making reflectors 
for automobiles, mirrors, radio tubes, and thermos bottles. 
Dextrose reduces an alkaline solution of copper sulfate and 
sodium potassium tartrate (or sodium citrate), known as 
Fehling’s solution. Thus, when this solution is boiled with 
dextrose (or any reducing sugar), a reddish, insoluble copper 
compound (cuprous oxide, CU 2 O) is formed. This method is 
often used as a test for reducing sugar. 

269. Lactose. — This sugar occurs in milk and is some¬ 
times called milk sugar. Cow’s milk contains from 3 to 5 
per cent of lactose. Lactose is not so sweet or soluble as 
sucrose. A solution of lactose reduces Fehling’s solution. 

Lactose is not fermented by ordinary yeast, but a special 
ferment, called lactic ferment, converts it into alcohol and 
lactic acid. The lactic acid gives milk its sour taste and also 
assists in curdling the milk, i.e., in changing the casein into 
a clot or curd. Lactose is obtained from whey. This is the 
liquid left after the solids have been pressed from milk which 
has been curdled by rennet in the manufacture of cheese. 
Lactose is used in infant foods and certain medicines. 

260. Maltose. — This sugar is formed from starch by 
malt. The transformation is caused by diastase, which is 
formed in the malt by allowing moist barley to sprout in a 


196 


A BRIEF COURSE IN CHEMISTRY 


warm place. Maltose is also formed from starch by the 
ptyalin in the saliva when food is well chewed. 

Maltose ferments readily with yeast, forming first dextrose 
and ultimately alcohol and carbon dioxide. Maltose re¬ 
duces Fehling’s solution. 

261. Starch. — This substance is found in wheat, corn, 
and all other grains, in potatoes, beans, peas, and similar 
vegetables, and also in rice, sago, tapioca, and nuts (espe¬ 
cially the chestnut). Many parts of plants contain starch, 
especially the root, seed, and fruit. 

The food value of vegetables depends largely on the starch 
they contain. Large quantities of starch are consumed as 
food, and in the manufacture of glucose and adhesives. 



Fig. -101. — Starch grains (magnified)—wheat (left), rice (center), corn 

(right). 

Starch, as usually seen, is a white, compact mass, but it 
really consists of granules which differ in size and shape with 
the plant (Fig. 101). 

Starch is not soluble in water. The granules are enveloped 
in an insoluble membrane of cellulose (§ 396). But if the 
starch is boiled with water, the membrane softens, the 
swollen grains burst through the membrane and form a clear 
liquid with the water. If a little water is used, a jelly-like 
mass is produced — the familiar starch paste. With cold 
water, starch forms an ordinary suspension. Whereas with 
considerable hot water, a colloidal suspension is produced, i.e., 
the particles are exceedingly small and do not settle (§ 95). 

Starch gives a blue colored substance when added to dilute 
iodine solution. The presence of starch in many vegetables 
and foods can be readily shown by grinding the substance 
in a mortar with cold water and adding a drop of dilute iodine 
solution. Starch does not reduce Fehling’s solution. 


SUGAR — STARCH — ACETIC ACID 


197 


Starch is a complex compound and its composition corre¬ 
sponds to the formula (C 6 Hio 05 )x- Starch is readily trans¬ 
formed into maltose by the ptyalin in the saliva. With 
dilute acids starch undergoes hydrolysis, thus: — 

(CeHioOs)^ -\- xli 20 = X ( CeKuOe ) 

Starch Water Dextrose 


262. Making bread. — Wheat flour contains about 70 
per cent of starch. The remainder is chiefly water and 
gluten. In making bread, the flour, water, fat, sugar, and 
yeast are thoroughly mixed into dough, which is put in a 
warm place to rise. Fermentation begins at once. Fer¬ 
mentation is the conversion of an organic compound, like 
starch or some sugars, into simpler compounds by the action 
of enzymes, e.g., diastase and 
ptyalin. 

By fermentation in the 
dough, alcohol and carbon di¬ 
oxide are formed. The gases 
escape in part through the 
dough, which becomes light and 
porous. When the dough is 
baked, the heat stops the action 
of the enzymes; but the alco¬ 
hol, carbon dioxide, and some 
water escape and puff up the 
mass still more. The heat, 
however, soon hardens the 
starch, gluten, etc., into a firm, 
porous loaf. 

263. Destructive distillation 
of wood. — Wood is composed 
essentially of cellulose (§ 396) 
and related organic compounds, 

water, and mineral matter. Cellulose is a compound of 
carbon, hydrogen, and oxygen. When wood is heated in 
large iron furnaces from which air is excluded, the cellulose 
and the other organic compounds decompose. In the labo¬ 
ratory a simple apparatus may be used (Fig. 102). Volatile 



Laboratory apparatus 
for the destructive distillation of 
wood. Wood is put in A and 
heated. The volatile products 
condense in C, charcoal is left in 
A, and gaseous products may be 
burned at E. 




















198 


A BRIEF COURSE IN CHEMISTRY 


products distill over with the water, and about thirty per 
cent of the wood remains in the furnace as wood charcoal 
(§ 43). This process is called dry or destructive distillation 
to distinguish it from ordinary distillation (§ 82). The dis¬ 
tillate is a brown liquid called pyroligneous acid. It contains 
acetic acid (about 10 per cent) and methanol. 

264. Acetic acid. — The acetic acid (CH3COOH) in the 
wood distillate is obtained by neutralizing the distillate with 
lime, decomposing the calcium acetate thereby formed, and 
distilling off the liberated acetic acid. 

Ordinary commercial acetic acid contains about thirty 
per cent of the acid, whereas glacial acetic acid is very con¬ 
centrated — nearly 100 per cent acid. 

Vinegar contains from 4 to 6 per cent of acetic acid. It is 
prepared by the transformation of sugars into ethyl alcohol, 
which is then changed into acetic acid by oxidation (through 
the action of bacteria). Substances containing fermentable 
sugars (§ 256), e.gf., fruit juices, cider, and molasses, ferment 
slowly when exposed to the air, which always contains bac¬ 
teria, and the ethyl alcohol becomes acetic acid, thus: — 

CeHisOe —^ C2H5OH -^CHsCOOH 

Fermentable Sugar Ethyl Alcohol Acetic Acid 

Acetic acid is a weak acid (§ 185). It forms many useful 
salts, e.g., sodium acetate (CHsCOONa), lead acetate 
((CH 3 COO) 2 Pb), and Paris green. The last named is a 
complex compound containing arsenic and copper and is used 
to kill potato bugs and other insect pests (§415). 

A test for acetic acid is the formation of ethyl acetate by 
interaction with ethyl alcohol. If a few drops of concentrated 
sulfuric acid are added to a mixture of acetic acid and ethyl 
alcohol, and the whole warmed gently, a pleasant fruitlike 
odor is detected. This is due to ethyl acetate which belongs 
to a class of compounds called esters (Pt. II, § 395). 

265. Methanol. — This is an alcohol and was formerly 
called “ wood alcohol.’’ It is used as a fuel, being the alcohol 
in “ solid alcohol ” (not solidified alcohol, but a mixture of 
methanol and a non-combustible gelatinous substance). 
Large quantities are used as a solvent for shellac gum; the 


SUGAR — STARCH — ACETIC ACID 199 

odor of the liquid called “shellac’’ is due to methanol. 
Some methanol is used to prepare denatured alcohol (§ 266). 
Methanol is a dangerous poison. 

Methanol is now largely manufactured by a synthetic 
process, which is much cheaper than the older process of 
distilling wood. A mixture of carbon monoxide and hydro¬ 
gen at a high pressure and temperature is passed over a suit¬ 
able catalyst. The equation is: — 

CO + 2H2 = CH3OH 

Carbon Monoxide Hydrogen Methanol 


266. Alcohols. — These are compounds of carbon, hy¬ 
drogen, and oxygen. There are many alcohols. One — 
methanol or methyl alcohol — has just been described (§ 265). 
The other important one is ethyl alcohol (sometimes called 
by the old name grain alcohol or the new name ethanol). 
Ethyl alcohol has the formula C2H5OH. 

Ordinary commercial ethyl alcohol contains 4 to 5 per cent 
of water. Absolute alcohol is 100 per cent ethyl alcohol. 
Denatured alcohol is essentially a mixture of 100 parts ethyl 
alcohol, 10 parts methanol, and a small proportion of some 
poisonous, or unpalatable, substance, such as benzene, pyri¬ 
dine, or kerosene. (There are many legal recipes for prepar¬ 
ing specially denatured alcohol.) Denatured alcohol is not 
suitable for use as a beverage, but specially denatured kinds 
can be used for industrial processes. 

Ethyl alcohol is used in manufacturing varnishes, celluloid, 
collodion, rayon (artificial silk), extracts, perfumes, dyes, 
ether ((C2H5)20), chloroform (CHCI3), iodoform (CHI3), 
and numberless other organic compounds. It is indispen¬ 
sable in many industries. 

Ethyl alcohol is manufactured either by the fermentation of the sugar 
left in the molasses, which remains from the extraction of crystallizable 
sugar in the sugar cane, or from vegetables and grains containing much 
starch, e.gr., potatoes and corn. If starch is used, it must be first changed 
into a fermentable sugar by treatment with an enzyme called diastase 
which is obtained from sprouting barley (malt). Yeast is added to the 
molasses or the sugar and an enzyme (zymase) in the yeast converts 
the sugar into alcohol and carbon dioxide. An equation is: — 


200 


A BRIEF COURSE IN CHEMISTRY 


C 6 H 12 O 6 = 2 C 2 H 6 OH + 2 CO 2 

Dextrose Alcohol Carbon Dioxide 

A test for ethyl alcohol is the formation of ethyl acetate 
(see end of § 264 ). 


EXERCISES 

1. State the properties of sucrose. 

2 . Compare dextrose and levulose. How are these sugars related 
to sucrose? 

3. What is the test for dextrose and similar sugars? 

4. Define and illustrate by means of sugars (a) hydrolysis and (b) fer¬ 
mentation. 

6. How would you show that a leaf contains starch? 

6 . How is acetic acid made ? 

7. What is (a) methanol, (6) ethanol ? 

PROBLEMS 

(See Problems at the end of Chapters VIII and IX.) 


SUGGESTIONS FOR LABORATORY WORK 
(References are to NewelFs Laboratory Exercises in Chemistry) 

Exercise 54 — Sugars 

Exercise 55 — Starch. 

Exercise *49 — Distillation of Wood — T. 

Exercise S41 (a) — Esters (Test for Acetic Acid). 

Exercise S41 (a) —Esters (Test for Ethyl Alcohol). 

SUPPLEMENTARY SECTIONS FROM PART II 

393. Manufacture of soap. 

394. The cleansing action of soap. 

396. What is soap? 

396. Cellulose. 

397. Derivatives of cellulose. 

398. Paper. 

At this point selections may be made from the sections in Topics 
XVII (Food and Nutrition), XVIII (Phosphorus — Arsenic Insecticides 
— Alloys of Antimony and Bismuth), XIX (Arrangement of the Ele¬ 
ments by Atomic Weights and by Atomic Numbers), XX (Fluorine — 
Bromine — Iodine), XXI (Silicon Dioxide — SiHcates — SiHcon Car¬ 
bide — Silicon Tetrafluoride — Glass). 


CHAPTER XXI 

METALS IN GENERAL 

267. Metals and non-metals. — The elements are often 
divided into metals and non-metals. The division is based 
largely on the conspicuous physical properties of the elements. 
Thus, the opaque, lustrous, more or less heavy, hard, ductile, 
malleable, tenacious solids are called metals. Whereas, all 
elementary gases, the liquid bromine, and solids such as 
carbon, sulfur, phosphorus, and iodine are called non-metals. 

A few elements sometimes act as metals and at other times 
as non-metals. Aluminum and antimony belong to this 
border-line class. 

Table VII shows the metals and non-metals in their usual class. The 
elements that may be in both classes are inclosed in a parenthesis. 


TABLE VII. — Metals and Non-Metals 


Metals 

Non-Metals 

Sodium 

Magnesium 

_ 

Hydrogen 

Oxygen 

Potassium 

Zinc 

Chromium 

— 

Sulfur 

— 

— 

— 

Boron 

— 

Copper 

Mercury 

Manganese 

— 

Fluorine 

Silver 

— 

— 

Carbon 

Chlorine 

Gold 

Aluminum 

Iron 

Silicon 

Bromine 

— 

— 

Cobalt 

— 

Iodine 

Calcium 

Tin 

Nickel 

Nitrogen 

— 

Strontium 

Lead 

— 

Phosphorus 

(Aluminum) 

Barium 

(Arsenic) 

Platinum 

Arsenic 

— 


Antimony 


(Antimony) 

(Tin) 


Bismuth 



(Chromium) 

(Manganese) 




201 














202 


A BRIEF COURSE IN CHEMISTRY 


268. Occurrence of metals. — Only a few metals are 
found in the native, or free, state in the earth’s crust, the 
important ones being gold, silver, copper, and platinum. 

Most of the metals occur in the earth’s crust as compounds, 
which are often called minerals. And the minerals from 
which metals can be profitably extracted are called ores. 

The most abundant classes of metal compounds are 
oxides, sulfides, and carbonates; other classes are chlorides, 
sulfates, phosphates, and silicates. Many ores contain 
arsenic. 

269. Properties of metals. — All have a metallic luster, 
i.e., the marked property of refiecting light from their pol¬ 
ished or untarnished surfaces. All are opaque except in 
very thin films. 

The color of many is white (silvery), though the tint varies. 
Thus, silver, sodium, aluminum, mercury, magnesium, iron, 
and tin are nearly pure white; bismuth is reddish white. 
Copper is the only red metal, and gold the only yellow one, 
which are elements. 

Most metals are malleable and ductile, Yc., they may be 
hammered or rolled into sheets and drawn into wire. Gold, 
copper, silver, iron, platinum, and aluminum possess both 
these properties (malleability and ductility) to a marked 
degree; while lead, zinc, and tin are very malleable though 
not so ductile. Antimony and bismuth are brittle. 

The hardness of metals varies. At ordinary temperatures 
sodium and lead can be cut easily with a knife, and so on 
through the list up to iridium, which is as hard as steel. 

In specific gravity {i.e., weight compared with the weight 
of an equal volume of water), the metals range between lith¬ 
ium, which is about 0.5, and osmium 22.5. Sodium and 
potassium are fighter than water, while magnesium has the 
specific gravity 1.75, and aluminum 2.58. 

Metals are good conductors of heat and electricity. They 
vary in this property. Silver, copper, and aluminum are 
the best conductors, and have therefore many practical 
applications, especially copper. 

Several of the distinctive chemical properties of metals 
already have been studied. 


METALS IN GENERAL 203 

1. Many metals form oxides when exposed to air or heated in oxygen. 
Iron is a conspicuous example. 

2. Metallic oxides dissolve in water and produce bases, e . g ., calcium 
oxide (CaO) and water form calcium hydroxide (Ca(OH) 2 ). 

3. Certain metals react with water. Some, like sodium, react with 
cold water, others, hke calcium, with hot water, and still others, Uke 
iron, magnesium, and zinc, with steam. Hydrogen is always produced 
together with an oxide or hydroxide. 

4. Certain metals react with acids, f.e., they displace hydrogen from 
acids and produce salts. (Nitric acid behaves exceptionally, § 203 .) 
Not all metals displace hydrogen from acids; hence we have a displace¬ 
ment series divided by hydrogen (see § 72 and 6 below). 

5. Metals form cations — positively charged ions — when salts are 
dissolved in water (§ 171 ), and in the electrolysis of salts metallic ions 
migrate to the cathode where they become atoms which (1) are deposited 
as metals, e . g ., copper or (2) interact with water and form secondary 
products, e . g ., sodium forms hydrogen and sodium hydroxide (§ 181 ). 

6. Metals displace other metals, as well as hydrogen, from solutions. 
They vary in their displacing power and can be arranged in a displace¬ 
ment (or electromotive) series based on this power. Magnesium is at 
one end, gold at the other, and hydrogen about midway (§§ 72 , 462 ), 
thus: Magnesium, aluminum, zinc, iron, lead, hydrogen, copper, 
mercury, silver, gold. 

270. Extraction of metals. — The series of operations by 
which metals are extracted from their ores and prepared for 
manufacture into useful articles is called jnetallurgy. Typi¬ 
cal processes are fully described in the chapter on iron and 
in the topics devoted to copper, aluminum, zinc, and lead. 

271. Alloys. — The term metallurgy also includes the 
preparation of alloys. These are mixtures, or compounds, of 
two or more metals. Some fused metals mix in all propor¬ 
tions, while others seem to form definite compounds. Alloys 
look like metals, indeed certain alloys, e.g., some of gold, can 
hardly be distinguished from the pure metal. The proper¬ 
ties of alloys vary with the constituents and their propor¬ 
tions. Some alloys have many industrial uses. For exam¬ 
ple, iron alloys (§ 315) are indispensable in automobiles, 
crushing machinery, and tools; copper alloys, known as 
brass and bronze, are widely used; and the aluminum alloy 
called duralumin is made into articles which range from 
kitchen ware to airships. Alloys in which mercury is a 
constituent are called amalgams. 


204 


A BRIEF COURSE IN CHEMISTRY 


EXERCISES 

1. Define (a) metallic luster, (6) malleable, (c) ductile, (d) specific 
gravity. 

2. Name four distinctive (a) physical and (b) chemical properties of 
metals. 

3. Define and illustrate (a) mineral and (b) ore. 

4. What are (o) alloys, (6) amalgams ? Illustrate each. 

PROBLEMS 

(See Problems at the end of Chapters VIII and IX.) 

1. What is the specific gravity of gold, if a piece weighs 4.676 gm. in 
air, and loses 0.244 gm. when weighed in water? 

(Note : Specific gravity equals the weight in air divided by the loss 
of weight in water.) 

2. A piece of aluminum weighs 150 gm. in air and 75 gm. in water. 
What is its specific gravity ? 

3. A piece of iron weighs 292.8 gm. in air and 255.3 gm. in water. 
What is its specific gravity ? 


SUGGESTIONS FOR LABORATORY WORK 

(References are to Newell’s Laboratory Exercises in Chemistry) 

Exercise *61 — Displacement of Metals. 

Exercise 62 — Flame Tests for Metals. 

Exercise S56 — Tests for Metals. 

Exercise S57 — Tests with Borax Beads. 

Exercise S58 — Cobalt Nitrate Tests. 


CHAPTER XXII 


SODIUM AND ITS COMPOUNDS 



272. Manufacture of sodium. — Sodium is manufactured 
by the electrolysis of melted sodium hydroxide — a method 
discovered by the English 
chemist Davy (Fig. 103). 

Figure 104 is a sketch of 
one form of apparatus 
used at Niagara Falls, 
where many electrical in¬ 
dustries are located. 

The body of the steel cyh 
inder (S) rests within a heated 
flue (not shown). The cathode 
(C) is iron and the connected 
carbon rods (AA) constitute 
the anode. A cylindrical col¬ 
lecting pot P surrounds the end 
of the cathode, and being per¬ 
forated at the lower end per¬ 
mits the circulation of the 
fused hydroxide but prevents 
the escape of the melted sodium 
in the upper part. 

The sodium hydroxide in 
the neck B is solid, but is kept 
melted and at about 300° C. in 
the main part by auxiliary gas 
flames. 

The sodium hydroxide conducts an electric current just as a solution 
does. As the electrolysis proceeds, sodium and hydrogen ions (Na+ and 
H+) migrate to the cathode (C). Here they are discharged, thereby 
producing sodium and hydrogen, which rise, and collect in the cylindri¬ 
cal pot P. The hydrogen escapes to some extent from under the cover, 
but enough always remains in the upper part of P to protect the sodium 

205 


Fig. 103. — The English chemist Davy 
(1778-1829), who first prepared sodium 
by the electrolysis of sodium hydroxide 
in 1807. 



206 


A BRIEF COURSE IN CHEMISTRY 


from the air. The molten sodium, which floats on the top of the fused 
sodium hydroxide, is skimmed off at intervals and poured into molds. 

Oxygen ions (0“” from the OH ions) migrate to the anode, are dis¬ 
charged, and produce oxygen (gas), which escapes through the pipe O 
without coming in contact with the 
sodium or hydrogen. 


273. Physical properties of so¬ 
dium. — Sodium is a silver-white 
metal. It is so soft that it can 
be easily cut with a knife. It 
floats on water, since its speciflc 
gravity is only about 0.97. 

274. Chemical properties of so¬ 
dium. — Heated in air, sodium 
melts at 96° C. At a higher tem¬ 
perature it volatilizes and burns 
with a brilliant yellow flame, 
forming sodium peroxide (Na 202 ). 
This intense yellow color is char¬ 
acteristic of sodium and its com¬ 
pounds and is the usual test for 
sodium (free or combined). (Po¬ 
tassium compounds, which are much like the corresponding 
sodium ones, color the flame a pale violet — thus being 
readily distinguished from sodium compounds.) 

In moist air the bright surface of sodium quickly tarnishes. 
Hence the metal usually has a yellow or brownish coating 
(instead of a shiny surface). It is kept under kerosene or a 
liquid free from water. 

Sodium decomposes water at ordinary temperatures, 
thus:— 

2Na + 2 H 2 O = H 2 + 2NaOH 

Sodium Water Hydrogen Sodium Hydroxide 

If held in one place upon the water by Alter paper, enough 
heat is generated by the reaction to set Are to the hydrogen, 
which burns with a yellow flame,, owing to the presence of 
volatilized sodium (§ 69). It combines vigorously with 
many non-metals, especially chlorine. If chlorine is passed 
through a tube containing melted sodium, the two elements 



paratus for manufacturing 
sodium. 









































SODIUM AND ITS COMPOUNDS 


207 


combine with a brilliant flame, forming sodium chloride. 
In this way Davy in 1810 proved that common salt is sodium 
chloride. 

276. Sodium chloride. — This is the most important com¬ 
pound of sodium. It is familiar under the name of salt or 
common salt. It is one of the most abundant substances 
and is the chief source of sodium compounds. 

276. Preparation of common salt. — Salt is obtained from sea 
water, rock salt deposits, and brines. The sea water is evaporated, often 
by exposure to the sun, and the salt separates from the concentrated 
solution. Deposits of salt are found in many parts of the globe, the 
most important being in England and Germany. The salt is dug out and 
then purified by recrystallization from water. In some parts of the 
United States natural or artificial brines are evaporated in large vacuum 
pans and then crystallized to the proper size in special apparatus. 

The dampness of salt is due to traces of magnesium and calcium 
chlorides which absorb moisture from the air (§ 94). Pure salt does not 
absorb moisture. Sometimes certain harmless substances are added to 
salt to keep it dry. 

277. Properties and uses of sodium chloride. — Sodium 
chloride is rather uniformly soluble in water, 100 gm. of water 
dissolving about 36 gm. of salt at 20 ° C., and about 39 gm. 
at 100° C. (Fig. 42). It crystallizes in cubes, and does not 
contain water of hydration (§ 85). 

Salt is an essential ingredient of the food of man and ani¬ 
mals. Besides its universal domestic use, enormous quan¬ 
tities are used in making sodium carbonate, sodium hydrox¬ 
ide, chlorine (§ 121), and hydrochloric acid (§ 130). 

278. Sodium carbonate. — This substance (Na 2 C 03 ) was 
formerly obtained from the ashes of marine plants. Now 
it is manufactured by the Solvay process. 

279. The Solvay process. — This consists in saturating a 
cold concentrated solution of sodium chloride with ammonia 
and carbon dioxide. The equation is : — 

H 2 O + NaCl + NH 3 + CO 2 = NaHC 03 + NH 4 CI 

Water Sodium Ammonia Carbon Acid Sodium Ammonium 
Chloride Dioxide Carbonate Chloride 

The acid sodium carbonate is sparingly soluble in cold am¬ 
monium chloride solution, and hence is precipitated. The 

/ 


208 


A BRIEF COURSE IN CHEMISTRY 


acid sodium carbonate is filtered off and then changed into 
normal sodium carbonate by heating, thus: — 

2 NaHCOg = Na 2 C 03 + CO 2 + H 2 O 

Acid Sodium Sodium Carbon Water 

Carbonate Carbonate Dioxide 


In operating the Solvay process (Fig. 105), a concentrated solution 
of sodium chloride is saturated with ammonia (in D), warmed to 40° 

C., and run 


(by C, C, C) 
into a tower 
(B) fitted with 
perforated par¬ 
titions. Car¬ 
bon dioxide is 
forced into the 
bottom (at A), 
and forms car- 
bonic acid, 
which reacts 
with the am¬ 
monium hy¬ 
droxide to 
produce am¬ 
monium bicar¬ 
bonate, 
NH 4 HCO 3 . 
This salt re¬ 
acts with the 
sodium chlo¬ 
ride and pro- 
duces acid 
into E, F. 



Fig. 105. — Sketch of the apparatus for manufacturing 
sodium carbonate (and bicarbonate) by the Solvay pro¬ 
cess. (Pumps marked P.) 


sodium carbonate and ammonium chloride. Both 
The precipitate of acid sodium carbonate is filtered off and changed by 
heating into normal sodium carbonate (Na2C03). The ammonium chlo¬ 
ride solution is pumped into G, where it reacts with calcium hydroxide 
forced in through H; the liberated ammonia passes through I into D, 
and the calcium chloride is drawn off through J. The carbon dioxide 
liberated by heating the sodium bicarbonate is used in B. 


280. Properties and uses of sodium carbonate. — Crys¬ 
tallized sodium carbonate (Na 2 C 03 .10 H 2 O), often called 
soda crystals, sal soda, or washing soda, effloresces, i.e., it 
slowly loses its water of hydration when exposed to air (§86). 
When heated; it first dissolves in its water of hydration, and 












































































































SODIUM AND ITS COMPOUNDS 209 

finally changes into the white anhydrous salt (Na 2 C 03 ) 
which is called soda ash or calcined soda. 

It is soluble in water, and forms an alkaline solution which 
is widely used as a cleansing agent; hence the name wash¬ 
ing soda. 

Enormous quantities are consumed in the manufacture of 
glass, water glass (sodium silicate), soap, borax, sodium 
hydroxide, and many other useful substances. 

281. Hydrolysis of sodium carbonate. — The alkaline re¬ 

action of sodium carbonate is due to hydrolysis and can be ex¬ 
plained in terms of the ionic theory. Hitherto water has been 
called a non-electrolyte, i.e., it does not ionize. As a matter 
of fact it does form the ions H+ and OH“, but to such a very 
slight extent that they have little or no effect in most cases. 
Under certain conditions, however, these ions (H+ and OH“) 
participate in reactions, e.g,, with sodium carbonate. Sodium 
carbonate ionizes into 2 Na+ and CO 3 , but the unstable 

C 03 -ions form HC 03 -ions with the H-ions from the slightly 
dissociated water. This removal of H-ions finally leaves in 
the solution sufficient OH-ions to produce an alkaline reac¬ 
tion. Equations for these ionic reactions are: — 

H 2 O = H+ + OH- 
Na 2 C 03 = 2 Na+ + CO 3 -- 
2 H+ + CO 3 - - = H 2 CO 3 

We may define hydrolysis as a chemical change in which 
water is an essential factor. A more restricted definition is 
the interaction of the ions of water with the ions of certain 
salts. A salt derived from a weak acid {e.g., H2CO3) and a 
strong base {e.g., NaOH) gives an alkaline solution. On the 
other hand, a salt derived from a strong acid (e.g., H2SO4) 
and a weak base {e.g. , Cu (OH)2) gives an acid solution. Thus, 
copper sulfate solution has an acid reaction. 

282. Sodium bicarbonate. — This substance is prepared 
by the Solvay process (see above), or by treating sodium 
carbonate solution with carbon dioxide. 

Sodium bicarbonate is a white powder, less soluble in water 
than the normal sodium carbonate, When heated alone or 


210 


A BRIEF COURSE IN CHEMISTRY 


when mixed with an acid or an acid salt, sodium bicarbonate 
liberates carbon dioxide. This property early led to its use 
in cooking, and gave the names cooking soda, baking soda, 
or simply soda. 

Sodium bicarbonate is sometimes called acid sodium car¬ 
bonate (though it is nearly neutral to litmus) and also sodium 
hydrogen carbonate. The neutral reaction of a solution 
of sodium bicarbonate is due to the fact that neither of its 
ions (Na+ and HCOs”) affects litmus. It should be noted 
that the name “ acid ” sodium carbonate emphasizes the 
method of formation (from carbonic acid), not the properties 
of the salt (§ 169). 

283. Baking powders. — Sodium bicarbonate is an essen¬ 
tial ingredient of baking powders and of the various mixtures 
(except yeast) used to raise bread, cake, and other food. 
The other ingredient is a substance which has a weak acid 
reaction, such as acid calcium, or acid sodium, phosphate 
(CaH 4 (P 04)2 or NaH 2 P 04 ), cream of tartar (acid potassium 
tartrate (KHC 4 H 4 O 6 )), or alum (K 2 Al 2 (S 04 ) 4 ). Commercial 
baking powders contain a small proportion of starch or flour, 
which prevents (or retards) premature chemical action. 

If baking powder is mixed with water, carbon dioxide is 
slowly liberated. With a tartrate baking powder the equa¬ 
tion is: — 


KHC 4 H 4 O 6 + NaHCOs = CO 2 + NaKC 4 H 406 + H 2 O 

Cream of Sodium Carbon Sodium Potassium Water 

Tartar Bicarbonate Dioxide Tartrate 

When dough is raised with baking powder, or with a mixture 
of baking soda and cream of tartar, the escaping carbon 
dioxide puffs it up. Hence baking soda is often called 
saleratus — the salt that aerates (from the Latin words 
sal, salt and aer, air). 

284. Sodium hydroxide or caustic soda. — This substance 
is a white, crystalline-, brittle, corrosive solid. It absorbs 
water {i.e., deliquesces— §94). It also absorbs carbon 
dioxide rapidly from the air, and is thereby changed into 
sodium carbonate. It is very soluble in water, and dis¬ 
solves with rise of temperature. The solution is strongly 


SODIUM AND ITS COMPOUNDS 


211 


alkaline and disintegrates many organic substances, e.g., 
wool, but not cotton. Hence the term caustic. It is a 
strong base and its solution contains a high per cent of hy¬ 
droxyl ions (§ 185). The solution is sometimes called lye or 

soda lye. 

Large quantities are used in refining petroleum and vege¬ 
table oils, and in making soap, paper pulp, phenol (carbolic 
acid, CeHsOH), and chemicals such as chlorates, hypochlo¬ 
rites, and nitrites. 

285. Manufacture of sodium hydroxide. — The chemical 
process consists in boiling a dilute solution of sodium car¬ 
bonate with calcium hydroxide; the main change is rep¬ 
resented thus: — 

Ca(OH )2 + Na 2 C 03 = 2 NaOH + CaCOg 

Calcium Sodium Sodium Calcium 

Hydroxide Carbonate Hydroxide Carbonate 

The solution of sodium hydroxide is filtered from the 
insoluble calcium carbonate and evaporated. The residue 
is heated, and the molten mass is then poured into small 
cylindrical molds about the diameter of a lead pencil or into 
large iron barrels called drums. Some is made into flakes. 

In the electrolytic process, a solution of sodium chloride 
is used. When an electric current is passed through such a 
solution, sodium hydroxide and hydrogen are produced at 
the cathode and chlorine at the anode. The cells are con¬ 
structed to prevent secondary action between the sodium 
hydroxide and chlorine (which would produce sodium hypo¬ 
chlorite or bleach liquor). A view of the cell room in an 
electrolytic alkali plant is shown in Fig, 106. 

Several types of electrolytic apparatus are used. Only one — the 
porous diaphragm type — will be described here. An example of this 
type is shown in Fig. 107. The liquid can penetrate the porous dia¬ 
phragm. Hence the diaphragm does not interfere with the flow of the 
electric current, but it does prevent the mixing of the two solutions. The 
porous diaphragm is a sheet of asbestos (mixed with iron oxide), which 
is supported on the perforated iron cathode. The graphite anode dips 
into the sodium chloride solution in the middle compartment; here the 
solution is kept at a certain level by regulation of the inflow. The 
outer compartment contains kerosene. 


212 


A BRIEF COURSE IN CHEMISTRY 



Fig. 106. — The cell room of a plant for making sodium hydroxide by elec¬ 
trolysis. 


When the current is passing, the sodium ions (Na+) migrate to the 
cathode, lose their charges, and become sodium atoms (Na). The 
sodium atoms interact with the water and form sodium hydroxide 
and hydrogen. The sodium hydroxide drops through the kerosene 

to the bottom of the outer compart¬ 
ment, and is drawn off through C; the 
hydrogen escapes through B. Simi¬ 
larly, chlorine ions (Cl~) migrate to the 
anode, lose their charges, and become 
chlorine atoms (Cl) which unite and 
escape as chlorine gas (CI 2 ) through A. 
The equations may be written thus: 
(1) NaCl = Na+ + CU; (2) 2 Na+ = 
2 Na and 2 Cl" = Ch ; (3) 2 Na + 
2 H 2 O = 2 NaOH + H 2 . 

The sodium hydroxide solu¬ 
tion is treated (without filtering) 
as described above (second para¬ 
graph, this section). The chlo¬ 
rine is stored or compressed 
into metal cylinders or is made 
directly into bleaching mixtures. 
286. Sodium tetraborate or borax. — This substance is a 
white crystalline solid, having five or ten molecules of water 
of hydration. A common household form of borax is the 



Fig. 107. — Sketch of the dia¬ 
phragm type of apparatus for 
the manufacture of sodium 
hydroxide by the electrolysis 
of sodium chloride solution. 



























SODIUM AND ITS COMPOUNDS 


213 


powdered crystals. The crystals readily effloresce and crum¬ 
ble in the air (§ 186.) When heated, crystallized borax loses 
its water of hydration and swells into a white porous mass, 
which finally melts into a glass-like solid. This glassy borax 
dissolves metallic oxides and is colored by them. 

When borax is melted on the end of a looped wire, the 
transparent globule is called a borax bead. These beads 
assume different colors if the bead is touched with a minute 
fragment of a metallic compound and then heated in an 
oxidizing or a reducing flame (Fig. 108). The colors are 
characteristic of the metals. Thus, a copper bead is made 
blue-green by an oxidizing flame and red by a reducing flame. 



Fig. 108. — Testing for a metal with a borax bead in the oxidizing flame 
(left) and reducing flame (right). 

Borax beads are used in testing for metals, e.g., often in a 
small quantity of a mineral. 

Its power to react with oxides adapts borax for use in 
soldering and welding metals. Solder adheres only to clean 
metals, so a little borax is used to dissolve the film of oxide 
on the surfaces to be joined. 

Borax is chiefly used in the manufacture of enamels for 
coating iron ware. The so-called “ granite or “ agate ” 
ware and “ porcelain-lined ” vessels are made of iron coated 
with an easily fused borax glass called enamel. Considerable 
borax is used for preserving meat, fish, cheese, and other 
foods, because it prevents the growth of certain bacteria. 

A solution of borax has a slight alkaline reaction owing to 
hydrolysis (§ 281) ; hence it is sometimes used instead of 
soap as a cleansing agent. Some soaps contain borax. 

287. Sodium nitrate. — This substance is found abun¬ 
dantly only in Chile and is often called Chile saltpeter. It 













214 


A BRIEF COURSE IN CHEMISTRY 



is extracted by dissolving out the saltpeter from the earthy 
matter and crystallizing the solution. The liquid left from 

the crystallization is a 
source of iodine (§§ 436, 
436). It is a white or 
brownish solid, which be¬ 
comes moist in the air, 
owing to slight deliques¬ 
cence (§ 94). Large 
quantities are used as a 
fertilizer, either alone or 
mixed with compounds of 
potassium and of phos¬ 
phorus (§ 143). It is 
used in making nitric acid 
(§ 197) and potassium 
nitrate. 

288. Fertilizers con¬ 
taining potassium. — Po¬ 
tassium, like nitrogen and 
phosphorus, is essential 
to the life of plants and 
animals. Potassium salts 
are taken from the soil 
by plants and must be 
returned if the soil is to 
be productive. Sometimes wood ashes are applied to the 
soil. Usually the potassium salts, e.g., potassium chloride, 
are supplied in the form of fertilizer. Experiments show 
that many soils need potassium salts as plant food (Fig. 
109). (Compare §§ 143, 300.) 


Fig. 109. — Result of an experiment with 
potassium salts. The pot on the left 
received a complete fertilizer, while the 
pot on the right differed only in receiv¬ 
ing a fertilizer without potassium salts. 


EXERCISES 

1. Describe the manufacture of sodium. 

2. Summarize (a) the physical properties and (6) the chemical prop¬ 
erties of sodium. 

3. Give an outline of the manufacture of sodium carbonate. Write 
the essential equations. 

4. What is (a) soda, (6) soda ash, (c) sodium carbonate, (d) soda 




SODIUM AND ITS COMPOUNDS 


215 


crystals, (e) sal soda, (/) washing soda, (g) calcined soda, (h) acid 
sodium carbonate, (z) saleratus, (j) baking soda, (k) caustic soda? 

6 . Describe the manufacture of sodium hydroxide by (a) the chemi¬ 
cal process and (6) the electrolytic process. 

6 . What is the flame test for (a) sodium, (h) sodium compounds, 
(c) potassium compounds? 

7. Why does sodium carbonate form an alkaline solution? 

8. Give the name and formula of the ions in dilute solutions of 
(a) sodium hydroxide, (b) sodium chloride, and (c) sodium nitrate. 

PROBLEMS 

1 . Calculate the weight of sodium in (a) 20 gm. of NaOH, and (b) 20 
gm. of sodium bicarbonate. 

2. How many kg. of sodium hydroxide can be made from 1000 kg. 
of NaaCOa? 

3 . What volume of hydrogen is liberated by the reaction of 1 gm. of 
sodium on water? 

4 . What is the per cent of sodium in Na 2 S 04 ? 

6. Calculate the simplest formula corresponding to Na = 32.39 per 
cent, S = 22.54, O = 45.07. 


SUGGESTIONS FOR LABORATORY WORK 

(References are to Newell’s Laboratory Exercises in Chemistry) 

Exercise S51 — Sodium Bicarbonate — T. 

Exercise S52 — Sodium Chloride — T. 

Exercise *11 — Reaction between Sodium and Water — T. 
Exercise S15 B — Deliquescence — T. 


CHAPTER XXIII 


CALCIUM CARBONATE — OXIDE — HYDROXIDE — 
SULFATE — PHOSPHATE 

289. Calcium carbonate. — An abundant form is lime¬ 
stone; large deposits are found in many places. Limestone 
is a white or gray compact solid, but impurities, especially 
organic matter and iron compounds, produce many colored 



Courtesy Thomsen-Ellis Co. 
Fig. 110. — A marble quarry in Georgia. 


varieties. Much limestone contains sand and clay. Some 
varieties contain the fossil remains of plants and animals. 
Hard, compact, crystalline limestone is called marble ; it is 
extensively used as a building and an ornamental stone 
(Figs. 110, 111, 112). Calcite is the purest form of crystal¬ 
lized calcium carbonate. 

Enormous quantities of the different varieties of calcium 
216 



CALCIUM CARBONATE — OXIDE — HYDROXIDE 217 



Courtesy Thomsen-ElUs Co. 

Fig. 111. — Cutting the flutes in a marble column. 



carbonate are used in making lime, cement, glass (§ 449), 
and sodium carbonate (§ 279), and as a flux in the 
smelting of metals, e.g., 
iron (§ 350). 


290. Deposition of calcium 
carbonate. — Calcium carbon¬ 
ate is only very slightly solu¬ 
ble in water. But if the water 
contains carbon dioxide, the 
carbonate dissolves, owing to 
its transformation into the 
soluble acid calcium carbonate 
(CaH2(C03)2). 

Some underground water 
contains carbon dioxide, and 
as this water works its way 
along in limestone regions, 
the limestone is dissolved and 
caves are often formed or 
enlarged. When the water 
enters a cave and drips from 
the top, the water evaporates, 
or the gas escapes, or both, 
and calcium carbonate is re- Fig. 112. — Sawing a block of marble 
deposited, often forming sta- into slabs with a gang saw. 












218 


A BRIEF COURSE IN CHEMISTRY 


lactites and stalagmites. Stalactites hang from the roof like icicles, and 
are often exquisitely shaped. Stalagmites grow up from the floor. 
Sometimes the two formations meet and produce a column (Fig. 113). 
Mexican onyx is a variety of stalagmite. Vast deposits of this beau¬ 
tiful mineral are found in Mexico and Algeria. It is translucent and 
dehcately colored, and is used as an ornamental stone, especially for 

altars, table tops, mantles, 
soda fountains, and lamp 
standards. Travertine is an¬ 
other variety; it occurs near 
many springs in Italy. When 
fresh it is soft and porous, but 
it soon hardens and becomes 
a durable building stone in 
dry climates. A portion of 
the walls of the Colosseum 
and St. Peter’s is travertine. 
Calcium carbonate which 
is dissolved in the ocean is 
extracted by marine organ¬ 
isms and transformed into 
shells and bony skeletons. 
The hard parts of these ani¬ 
mals accumulate in the mud 
on the ocean bottom, and 
subsequently form a part of 
the land. Chalk is the re¬ 
mains of shells of minute 
marine animals. When ex¬ 
amined under a microscope, 
a good specimen is seen to 
consist almost entirely of tiny shells (Fig. 114). Coral is calcium car¬ 
bonate and the vast accumulations in the sea are the skeletons of the 
coral animals. 

291. Chemical properties of calcium carbonate. — Cal¬ 
cium carbonate is decomposed by heat into lime (CaO) and 
carbon dioxide (§ 293). It interacts with acids (§ 55), e.g .: — 

CaCOa + 2 HCl = CaCb + H 2 O + CO 2 

Calcium carbonate itself is precipitated by the interaction 
of a soluble calcium salt and a soluble carbonate, thus: — 

CaCl 2 + Na 2 C 03 = CaCOs + 2 NaCl 

Calcium Sodium Calcium Sodium 

Chloride Carbonate Carbonate Chloride 






CALCIUM CARBONATE — OXIDE — HYDROXIDE 219 



This reaction may be used as a test for a soluble carbonate, 
or in preparing a very fine precipitate of calcium carbonate. 
The test for an insoluble carbonate is 
the liberation of carbon dioxide upon 
the addition of dilute hydrochloric 
acid. The liberated gas turns calcium 
hydroxide turbid. A simple way to 
make this .test is shown in Fig. 115. 

Purified calcium carbonate is called 
precipitated chalk, and is used exten¬ 
sively as the polishing ingredient of 
tooth powder. An impure variety is 
whiting ; and a mixture of whiting and 
linseed oil is putty. 

292. Lime. — This familiar sub¬ 
stance is calcium oxide (CaO). It is a 
hard, more or less porous, white solid. 

Lime when exposed to the air be¬ 
comes air slaked,” Le., it slowly ab¬ 
sorbs water and carbon dioxide, swells 
considerably, and soon crumbles to a 



Fig. 115. — Testing for an 
insoluble carbonate. The 
carbon dioxide liberated 
by dilute acid interacts 
with the drop of calcium 
hydroxide on the end of 
the glass tube and forms 
white, insoluble calcium 
carbonate. 















220 


A BRIEF COURSE IN CHEMISTRY 


powder. This powder is a mixture of calcium hydroxide and 
calcium carbonate; such a mixture is not suitable for many 
of the uses of lime. (See, however, hydrated lime, § 294.) 
Lime that is not air slaked is called quicklime or caustic lime. 

When just the right amount of water is slowly added to 
lime, the two substances combine and form calcium hydroxide 

If an excess of water is added 
quickly, the lime and water com¬ 
bine vigorously ; considerable 
heat is liberated, as is seen 
when mortar is being prepared. 
This operation is called slak¬ 
ing, and the product is slaked 
lime. Sometimes water leaks 
into barrels, cars, or buildings 
in which lime is stored, and the 
heat evolved causes a fire. 

Large quantities of lime are 
used in preparing mortar and 
plaster. Many useful chemi¬ 
cals are made from lime, e.g., 
bleaching powder (§ 125), cal¬ 
cium carbide (§ 50), sodium 
hydroxide (§ 285), calcium bi- . 
sulfite, and lime-sulfur mixtures 
(§ 204). Considerable lime is 
used in such industrial opera¬ 
tions as purifying illuminating 
gas, refining sugar, softening 
water, removing hair from hides, bleaching cotton cloth, and 
extracting metals from ores. In agriculture lime is added to 
sour soil to neutralize the acids. 

293. Manufacture of lime. — Lime is manufactured by heating lime¬ 
stone. The decomposition takes place according to the equation: — 

CaCOs = CaO + CO 2 

Calcium Carbonate Calcium Oxide Carbon Dioxide 

The carbon dioxide escapes; the lime is left. 

Limestone was formerly “ burned ” in a cavity on a hillside, and in 
some regions it is so prepared to-day. An arch of limestone is built 


or hydrated lime (§ 294). 



Fig. 116. — Sketch of a contin¬ 
uous vertical limekiln. 
































CALCIUM CARBONATE — OXIDE — HYDROXIDE 221 


across the cavity above the fire pit, and limestone is introduced until 
the kiln is full. The arch kilns have been largely replaced by contin¬ 
uous kilns, i.e., either rotary or vertical furnaces (Figs. 116, 117, 118). 

A sketch of a vertical kiln is shown in Fig. 116. The heat is produced 
at A, A. The hot air and gaseous products of combustion pass up 
through the lumps of limestone (fed in at the top) and heat it. The 
rising gases sweep out the carbon dioxide. A proper temperature (about 
750° C.) is necessary. The lime drops down through the furnace and 
is removed (at B, B). 

Modern vertical kilns are 60 feet or more high (Fig. 118). They 
operate continuously and produce from 25 to 60 tons a day according 



Fig. 117. — A rotary kiln for making lime from marble refuse. 


to size. Several towers are often connected with a source of heat, e.g., 
a gas producer or a furnace burning a smokeless fuel such as coke. 

294. Calcium hydroxide. — This substance is a white 
solid. It is manufactured by adding carefully just enough 
water to calcium oxide to produce the hydroxide. This 
hydrated lime, as it is called, is a fine white powder. If 
properly packed, it will keep indefinitely. It can be stored 
without danger of causing fire, and is suitable for the same 
purposes as the lime slaked just before use. 

Calcium hydroxide is sparingly soluble in water. It is un¬ 
like most solids in being more soluble in cold than in warm 
water. The solution has a bitter taste and a mild alkaline 







222 


A BRIEF COURSE IN CHEMISTRY 


reaction; it is called limewater. Exposed to the air, lime- 
water becomes covered with a thin crust of calcium carbonate, 
owing to interaction with carbon dioxide. For the same 
reason, limewater becomes milky or cloudy when carbon 
dioxide is passed into it. The formation of calcium carbonate 
in this way is the test for carbon dioxide. 

Limewater is prepared by carefully adding lime to consider¬ 
able water, allowing the mixture to stand in a stoppered 
bottle until the solid has settled, and then removing the clear 

liquid with a siphon. 
When considerable cal¬ 
cium hydroxide is sus¬ 
pended in the liquid, the 
mixture is called milk 
of lime. Ordinary white¬ 
wash is thin milk of lime. 

295. Mortar. — This is 
made by mixing lime with 
three or four times its 
bulk of sand, and then 
adding enough water to 
produce a thick paste. It 
slowly hardens or sets, 
owing to the evaporation 
of water and to the inter¬ 
action with carbon diox¬ 
ide. When placed between bricks or stones it holds them 
firmly in place, and is used to construct buildings, walls, 
foundations, etc. 

The sand gives bulk and rigidity; it also makes the mass 
porous and thus facilitates the change of the hydroxide to 
the carbonate. We can readily prove that a carbonate is 
formed by adding hydrochloric acid to a lump of old mortar ; 
the gas which is liberated is carbon dioxide. Hair is some¬ 
times added to make the mortar stick better, especially 
when it is used as plaster for walls. Hair is not necessary 
if cement is mixed with the mortar. 

296. Cement. — This is a kind of strong, firm mortar. 
Like ordinary mortar, it hardens in the air; unlike mortar. 



Fig. 118 . — A modern vertical limekiln 
in operation. 






CALCIUM CARBONATE — OXIDE — HYDROXIDE 223 


however, it hardens under water. The hardening takes place 
without carbon dioxide. 

The chemical changes which occur in the setting of cement 
are complex and not well understood. The products set 
into a compact, almost impervious solid. 

Cement is next to iron and steel in importance as a build¬ 
ing material. Immense quantities (about 100,000,000 barrels, 
each containing 380 lb.) are made annually in the United 
States. It is used in a great variety of structures — founda- 



Fig. 119. — A huge arch of cement in a railway bridge. 

tions, dams, bridges, tunnels, levees, fireproof buildings 
{e.g., garages), storage tanks, warehouses, grain elevators, 
floors, walks, and roads (Fig. 119). 

A mixture of cement, sand, water, and crushed stone (or 
coarse gravel) is concrete. This mixture is used, usually 
instead of cement alone, as construction material, particu¬ 
larly for foundations or for walls designed to withstand great 
pressure, e.g., bridges (Fig. 119). Sometimes concrete is 
reenforced by imbedding twisted steel rods or steel sheets 
in it; the concrete and steel form a firm union, if properly 
constructed. 












224 


A BRIEF COURSE IN CHEMISTRY 


Cement is manufactured from a mixture of limestone, clay, 
and sand. The raw materials must contain the proper 
proportions of the essential ingredients. The materials are 
ground very fine, thoroughly mixed, and heated in a steel 



Fig. 120. — Rotary kilns for making cement. 


furnace from 70 to 150 feet long and 6 to 8 feet in diameter, 
and lined with fire-brick. The furnace is inclined about 
15 degrees and rotates slowly (about once a minute); it is 
called a rotary kiln (Fig. 120). 

The process can be best understood by studying the simple sketch in 
Fig. 121. The powdered mixture is fed in at the upper end (C). As it 
gradually works its way along through the slowly rotating kiln it is 
heated to about 1500 C. by the flame and hot gases produced (inside 



the kiln) by burning oil or coal dust, which is forced in at A by a powerful 
air blast. The mixture forms a semifused, gray-black mass which drops 
out at B. The cooled lumps, called clinker, are mixed with about 2 
per cent of gypsum (calcium sulfate), and ground to an exceedingly fine 
powder. This powder, which is grayish, is cement. 















CALCIUM CARBONATE — OXIDE — HYDROXIDE 225 

297. Calcium sulfate. — Extensive deposits of different 
forms of calcium sulfate are found in many localities in the 
United States. Gypsum is the commonest form; it occurs 
as white masses which have the composition CaS 04.2 H 2 O. 
A translucent, crystallized variety of gypsum is called 
selenite. A variety of crystallized gypsum colored slightly 
by impurities {e.g., iron compounds) is called alabaster. 
The mineral anhydrite is anhydrous calcium sulfate (CaS 04 ). 

Gypsum is a rather soft solid. It is only slightly soluble 
in water. It is used as an ingredient of some fertilizers and 
in making plaster of Paris, paper, white paint, blackboard 
crayon, cement, and fireproof blocks. 

Plaster of Paris is made by heating gypsum to the proper 
temperature (about 145° C.). The equation is: — 

2 CaS04.2 H 2 O = (CaS04)2. H 2 O + 3 H 2 O 

If gypsum is moistened with water, it swells and quickly 
sets or solidifies to a hard mass which consists of a network 
of very small crystals. The equation for the setting of 
plaster of Paris may be written: — 

(CaS04)2. H 2 O + 3 H 2 O = 2 CaS04.2 H 2 O 

Plaster of Paris Water Crystallized Gypsum 

Plaster of Paris is used to coat plastered walls, to cement 
glass to metal, but more largely to make casts and repro¬ 
ductions of statues and small objects. Stucco is a mi^cture 
of glue and plaster of Paris. 

298. Calcium compounds and hardness of water. — Cal¬ 
cium sulfate is'slightly soluble in water, and calcium carbon¬ 
ate, as already seen, is changed into soluble acid carbonate 
(Ca(HC 03 ) 2 ) by water containing carbon dioxide. Water 
containing the sulfate and acid carbonate (or the correspond¬ 
ing magnesium salts) is called hard water. And water from 
which they are absent is often called soft water. 

Soap does not dissolve readily in hard water, but forms, 
at first, sticky, insoluble compounds with calcium (and mag¬ 
nesium) salts; hence a large quantity of soap must be used. 

Hard water, if used in boilers, forms a hard, clinging 
deposit or scale ” on the inside of the boiler tubes, thus 


226 


A BRIEF COURSE IN CHEMISTRY 


causing waste of heat (Fig. 122). A deposit one sixteenth 
of an inch thick causes a loss of about 20 per cent and one 
fourth inch nearly 50 per cent. 

Hardness due to acid calcium carbonate (or acid magne¬ 
sium carbonate) is called temporary hardness, because it can 
be removed by boiling. Temporary hardness can also be 
removed by adding the correct amount of calcium oxide or 
hydroxide to the water to change the soluble to the insoluble 
carbonate. 

Hardness due to calcium sulfate or chloride is called 
permanent hardness, because these salts can not be removed 
by boiling. (Magnesium sulfate and chloride, 
like the corresponding calcium salts, produce 
permanent hardness.) Permanently hard 
water can be softened by adding the neces¬ 
sary amount of sodium carbonate or sodium 
phosphate, which converts the calcium and 
magnesium salts into the insoluble carbonate 
or phosphate. This process is used on a large 
scale to soften boiler water. In the home, 
borax or ammonia may be used. 

Several industrial processes are used to 
soften boiler water. One uses a coarse sand¬ 
like substance called permutit (essentially an 
artificial sodium silico-aluminate), which in¬ 
teracts with calcium (and magnesium) compounds, and 
thereby removes them from the water. 

A sketch of the apparatus is shown in Fig. 123. The hard water 
enters at A and slowly filters through the upper layer of marble chips or 
clean gravel (R), the porous layer of the permutit (C), the lower layer of 
gravel (D), and out the soft water outlet (F). By this process the cal¬ 
cium (and magnesium) replaces the sodium, thus: — 

Calcium Sulfate Sodium Permutit = 

Calcium Permutit + Sodium Sulfate 

After about twelve hours, the calcium permutit accumulates to such 
an extent that the mixture no longer reacts. Then water containing 
sodium chloride is added (from 7) and allowed to remain long enough 
(about twelve hours) to regenerate the sodium permutit, thus: — 

Calcium Permutit -f 2 NaCl = Sodium Permutit + CaCl? 



Fig. 122. —Sec¬ 
tion of a boiler 
tube showing 
the scale de- 
posited by 
hard water. 


CALCIUM CARBONATE — OXIDE — HYDROXIDE 227 

The permutit is cleaned and the calcium chloride removed by forcing 
water up through F and out at G, and the filter is ready for use again. 

299. Calcium phosphate. — Calcium phosphate (Ca 3 (P 04 ) 2 ) 
is the calcium salt of phosphoric acid (H3PO4). It is abun¬ 
dant, being the chief ingredient of bones and phosphate rock. 
Bones are about 80 per cent calcium phosphate. Phosphate 
rock is the name given to the 
hardened remains of land and 
marine animals. It occurs in 
large beds in Florida, Tennessee, 
and North and South Carolina 
(Fig. 124). Its main constituent 
is calcium phosphate (Ca 3 (P 04 ) 2 ). 

Artificial phosphate fertilizers are 
made by treating phosphate rock 
with sulfuric acid. 

300. Relation of phosphorus 
to life. — Phosphorus is essential 
to the growth of plants and ani¬ 
mals. Plants absorb soluble 
phosphates from the soil and 
store up phosphorus compounds, 
especially in their seeds. Ani¬ 
mals eat this vegetable matter, 
assimilate the phosphorus com¬ 
pounds, and deposit them to 
some extent in the bones, as well 
as in brain and nerve tissue. 

Some of the complex phosphorus 
compounds in the food consumed 
by animals are transformed into 
phosphates. These are eliminated to a large extent, and 
thus often find their way back into the soil. Here they 
are taken up again by plants, converted into complex com¬ 
pounds, stored up in the seeds and other parts, which are in 
turn eaten by animals. And so the process goes on a 
phosphorus cycle analogous to the carbon cycle (§ 69, 
Fig. 24). 



the permutit process. 



































228 


A BRIEF COURSE IN CHEMISTRY 


301. Phosphate fertilizers. — In order to furnish plants 
with phosphorus compounds, various phosphorus-bearing 
substances are added to the soil in the form of phosphates, 
especially the fertilizer made from phosphate rock. 

The beneficial results of adding (complete) fertilizer to 
the soil are strikingly shown in Fig. 125. In the wheat 
field (top), the left was fertilized and the right was unfer¬ 
tilized ; the left yielded 17 bushels per acre but the right 
only 5. The contrast is more conspicuous in the sorghum 
field (bottom) where the unfertilized part (left) yielded a 



Fig. 124. — Mining phosphate rock in Florida. The powerful stream of 
water washes the phosphate rock down into a pit from which both 
water and rock are pumped to washers, where sand is removed and low 
grade rock rejected. 

product which gave only 66 gallons of sirup per acre, whereas 
the right yielded 140 gallons. From the cotton field (middle) 
about 1350 pounds of seed cotton were obtained from the 
heavily fertilized part and only about 700 pounds from the 
scantily fertilized part. 

302. Manufacture of phosphate fertilizer. — Tricalcium phosphate 
(Ca 3 (P 04 ) 2 ), as it is sometimes called, is insoluble in water, and must be 
changed into the soluble mono-calcium salt (H 4 Ca(P 04 ) 2 ), so it can 
be evenly distributed through the soil and easily taken up by plants. 
This soluble salt is sometimes called “ superphosphate of lime.” When 












Fig. 125. — Contrast of fertilized and unfertilized fields. 

229 

























230 


A BRIEF COURSE IN CHEMISTRY 


phosphate rock is treated with sulfuric acid, the changes involved may 
be written thus: — 


Ca3(P04)2 + 2 H2SO4 = H4Ca(P04)2 + 2 CaS04 

Tricalcium Sulfuric “Superphosphate Calcium 

Phosphate Acid of Lime” Sulfate 


Ca 3 (P 04)2 + 3 H2SO4 = 2 H3PO4 + 3 CaS 04 

Phosphoric Acid 

Ca3(P04)2 + H2SO4 = H2Ca2(P04)2 + CaS 04 

Dicalcium Phosphate 


The mixture of “ superphosphate ’’and calcium sulfate when ground 
and dried is ready for use as a phosphate fertilizer. Usually “ super¬ 
phosphate ” is mixed with compounds of nitrogen (§ 143) and of 
potassium (§ 288) to form a complete fertilizer. 

The law requires the dealer to state the analysis of the fertilizer on the 
bag or label. The per cent of phosphorus is usually stated as per cent 
of P 2 O 6 , which is popularly called “ phosphoric acid.” 


303. Tests for calcium. —A white solid called calcium 
oxalate (CaC204) is formed by the interaction of ammonium 
oxalate and a dissolved calcium compound; it is insoluble 
in acetic acid but soluble in hydrochloric acid. Its forma¬ 
tion and properties serve as a test for calcium. The orange- 
red color imparted to the Bunsen flame is another test for 
calcium. (Strontium salts and barium salts, which resemble 
the corresponding calcium salts, produce a different colored 
flame — strontium salts crimson and barium salts green.) 


EXERCISES 

1. State the properties and uses of lime. How is it made? 

2. Describe the manufacture of cement. 

3. What is hard water? How does it act with soap? What is 
(a) temporary hardness and (6) permanent hardness ? How can each be 
removed? What is soft water? Why is rain water often called soft 
water ? 

4. Write equations for the reactions necessary to prepare (a) calcium 
nitrate from calcium carbonate, (6) calcium hydroxide from calcium 
chloride, (c) calcium carbonate from calcium hydroxide. 

6. State the tests for calcium. 

6. Express the following reactions by equations: (a) calcium hy¬ 
droxide and carbon dioxide form-and water; (6) calcium carbon¬ 
ate forms-and carbon dioxide. 




CALCIUM CARBONATE — OXIDE — HYDROXIDE 231 


PROBLEMS 

1. Calculate the simplest formula from the following per cents: 
Ca = 29.49, O = 46.92, S = 23.59. 

2. How many tons of hme can be made from 2000 tons of limestone 
(95 per cent pure) ? 

3. How many tons of water are needed to change 1500 tons of hme 
into calcium hydroxide? 

4. What per cent of calcium carbonate is calcium ? 


SUGGESTIONS FOR LABORATORY WORK 

(References are to Newell’s Laboratory Exercises in Chemistry) 

Exercise *57 — Calcium Carbonate, Oxide, and Hydroxide. 
Exercise *58 — Hardness of Water — T. 

Exercise S53 — Tests for Calcium in Compounds. 

Exercise S54 — Properties of Cement — T. 

Exercise S55 — Plaster of Paris — T. 


CHAPTER XXIV 


IRON AND STEEL —IRON COMPOUNDS 

304. Occurrence of iron. — Iron ranks fourth in abundance 
among the elements and second among the metals (§ 326). 



Fig. 126. — Sketch of a blast furnace 
and the process of smelting iron ore. 


Combined iron is found in 
most rocks, soils, and natural 
waters. It is assimilated by 
plants and animals and is 
essential to their life pro¬ 
cesses, being a constituent of 
chlorophyll (the green color¬ 
ing matter of plants) and of 
haemoglobin (the red coloring 
matter of blood). 

The important ores of iron are 
hematite (ferric oxide, Fe 203 ), 
limonite (hydrated ferric oxide, 
2 Fe 203 . 3 H 2 O), magnetite (mag¬ 
netic oxide, Fe 304 or Fe(Fe 02 ) 2 ), 
and siderite (ferrous carbonate, 
FeC 03 ). The most abundant ore 
and the chief source of iron and 
steel in the United States is hema¬ 
tite, which comes mainly from 
mines in the Lake Superior region. 

305. How iron is obtained 
from its ores. — Iron is ob¬ 
tained by reducing its oxides. 
The ore is mixed with a flux 
(usually limestone) and car- • 
bon (usually in the form of \ 
coke) and heated in a blast 1 
furnace. The carbon together I 


























IRON AND STEEL —IRON COMPOUNDS 233 


with carbon monoxide reduces the oxide to metallic iron. 
The flux converts the mineral impurities in the ore, called 
gangue {e.g., silicon and aluminum compounds), into fusible 
silicates called slag. The operation, which is called smelting, 
is carried out in a blast furnace (Figs. 126, 128). 

The blast furnace (Fig. 126) is a tower, about one hundred 
feet high and twenty feet in (inside) diameter at the largest 
part; it is narrower at the top and 
bottom than in the middle. It is 
built of steel and lined with fire 
brick, the walls being several feet 
thick. Near the bottom is a large 
pipe encircling the furnace and 
provided with outlet pipes, called 
tuyeres (T), through which large 
quantities of hot, dry air are forced 
(f.e., blasted) into the furnace. 

The air blast is preheated by 
being passed through chimney¬ 
like towers called stoves (Fig. 

127) which are between the pair 
of furnaces (Fig. 128). 

Another pipe, D, from six to 
ten feet in diameter (called the 
‘‘downcomer’’) permits the es¬ 
cape at the top of the large 

quantity of hot gases (50 per cent nitrogen and about 25 per 
cent each of carbon monoxide and dioxide). Gases from the 
“downcomer” D (Fig. 127) enter the stove at the bottom and 
meet air, at A. The burning gases pass up and down, heat 
the checkerwork (shown in cutaway part at left), and escape 
into the chimney at C. To operate, dry air is blown in at B 
which passes up and down the hot checkerwork, and out 
through E to the tuyeres T. 

When the blast furnace is in operation, charges of the 
proper proportions of ore, coke, and flux are hauled up an 
inclined track on the furnace in small cars, and introduced 
at intervals by dumping them upon the cone-shaped cover. 
The latter is constructed and operated so that the charge is 



Fig. 127. — Section of a blast 
furnace stove. 











































234 


A BRIEF COURSE IN CHEMISTRY 



distributed uniformly in layers of ore, coke, and limestone 
without letting out gases or interrupting the smelting. The 
hot, dry air which enters at the bottom, through the tuytos 
(T), changes most of the carbon (coke) in the lower part of 
the furnace into carbon dioxide and thereby generates 
intense heat. The carbon dioxide is largely reduced by the 
hot carbon (coke) above it to carbon monoxide, which rises 
through the furnace and reduces the iron oxide to iron. As 
the mixture of ore, coke, and limestone settles down through 


Fig. 128. — Blast furnaces (at ends) and stoves (in the middle). 

the furnace, smelting proceeds continuously, i.e., the ore is 
reduced to metal and the flux forms a slag (largely calcium 
silicate, CaSiOs). Both iron and slag sink, the molten iron 
finally falling through the slag to the bottom of the furnace, 
where they form two layers. 

The molten slag, which floats on the molten iron, is tapped 
off (at S), granulated by running it into water, and carted to 
the “ dump.’^ It is used to some extent as ballast or dust- 
preventer for railroads, and to make mineral wool, paving 
stones, and cement. 

The iron is tapped off (at I) into huge ladles, run from the 
furnace into molds of sand or iron (^.c., in a casting machine). 























IRON AND STEEL —IRON COMPOUNDS 235 


and allowed to solidify into blocks (called pigs ’0 weighing 
75 to 100 pounds. Such iron is called cast iron or pig iron. 
In some plants the molten iron is run into a casting machine 
(Fig. 129), or into huge vessels, called converters, and made 
into steel (§ 310). 

306. Cast iron. — The cast iron or pig iron that comes 
from a blast furnace is impure iron. It contains from 92 to 
94 per cent of iron, 2.5 to 4.5 of carbon, 1 to 3 of silicon, about 



Fig. 129. — Cast iron flowing from a ladle into one end of a long casting 
machine. 


0.7 each of manganese and phosphorus, and 0.02 to 0.05 of 
sulfur. 

The properties of cast iron depend on the rate at which the 
molten iron cools, as well as on the proportions of the ingre¬ 
dients other than iron. If the cooling takes place rapidly 
(as in the casting machine, Fig. 129), most of the carbon 
combines with the iron and forms a carbide (FesC), called 
cementite. This variety of cast iron is called white cast iron. 
But if the cooling occurs slowly (as in sand molds), much of 
the carbon remains uncombined as scales of graphite (called 
graphitic carbon). This variety of cast iron is known as 
gray cast iron. It is softer and less brittle than the white 
variety. It also melts at a lower temperature, forms a 
thinner liquid, and makes better castings. 




236 


A BRIEF COURSE IN CHEMISTRY 


Cast iron has a crystalline structure and is brittle; but 
it will withstand great pressure. It can not be welded nor 
forged, that is, hot pieces can not be united nor be shaped 
by hammering. But it can be cast, i.e., melted and formed 
into a desired shape by pouring the molten metal into a mold. 

Cast iron is the variety used in an ordinary iron foundry. 
Here the iron, which melts at about 1200° C., is heated in a 
furnace similar to a blast furnace, and when molten is poured 
into sand molds of the desired shape. Stoves, pipes, pillars, 
railings, radiators, parts of machines, and many other useful 
objects are made of cast iron. Considerable cast iron is made 
into steel (§ 309). 

Cast iron to which 5 to 20 per cent of manganese has been 
added is called spiegel iron, while ferro-manganese contains 
20 or more per cent of manganese; both are used in making 
steel. 

Cast iron is not attacked by alkalies and only slightly by 
concentrated acids. Concentrated sulfuric acid is trans¬ 
ported in iron tank cars. Cast iron interacts readily, 

however, with dilute acids. 

307. Wrought iron. — 
Wrought iron is made 
from cast iron by remov¬ 
ing most of the impurities 
(carbon, silicon, phos¬ 
phorus, and sulfur). This 
can be done by heating 
the cast iron with iron 
oxide. 

The process is conducted in a reverberatory furnace (Fig. 
130). The hearth B of the furnace is covered with iron ore 
(ferric oxide, Fe 203 ) and the charge of cast iron and flux is 
laid upon it. The long flame and hot gases from the fire on 
the grate G pass over the bridge E, are reflected down upon 
the charge by the sloping roof of the furnace, and melt the 
cast iron. Gases escape through 7. The charge, which 
rests on B, does not come in contact with the fuel on G. 
The carbon in the cast iron unites with the oxygen of the iron 
oxide and escapes as carbon monoxide. The silicon and 



Fiq. 130. — Reverberatory furnace. 
















IRON AND STEEL —IRON COMPOUNDS 237 


phosphorus are oxidized and react with the flux to form a 
slag; the manganese and sulfur (in the form of ferrous sul¬ 
fide) also become a part of the slag. The mixture is stirred 
or “ puddled ’’ with long rods, and as the impurities are re¬ 
moved, the mass becomes pasty. 

Finally, large balls, called blooms, are removed and ham¬ 
mered, or, more often, rolled between ponderous rollers. 
This operation squeezes out most of the slag. If the rolling 
is repeated, the quality of the iron is improved; the final 
rolling often leaves the iron in the desired commercial shape. 



Fig. 131. — Photomicrographs of wrought iron (left) and cast iron (right). 
The slag can be seen in the wrought iron. 


Wrought iron is the purest variety of commercial iron. It 
is practically pure iron, containing only 2 per cent (or less) 
of slag (Fig. 131). Wrought iron seldom contains more 
than 0.5 per cent of carbon and sometimes only 0.06 per cent, 
the average being about 0.15 per cent; the other elements 
are present in mere traces. 

If a specimen of iron is polished and then etched with weak 
acid, the microscope reveals the crystalline or other formation 
and the presence of carbon, slag, carbides, etc. Photographs 
of a treated specimen of wrought iron and cast iron as seen 
under the microscope are shown in Fig. 131. 

Wrought iron, unlike cast iron, is fibrous and can be bent. 
Since it softens at about 1000° C., it can be forged and welded. 
It is very malleable and ductile, can be readily rolled into 






238 


A BRIEF COURSE IN CHEMISTRY 


plates and sheets, and drawn into fine wire; in these forms 
the metal is very strong. 

Wrought iron rusts more rapidly than cast iron, and is 
also more vigorously attacked by acids and alkalies at a high 
temperature. It is not hardened by sudden cooling (§ 314). 

Wrought iron is made into wire, sheets, rods, nails, spikes, 
bolts, chains, anchors, and agricultural implements. It is 
less important than formerly, since it is being replaced by 
soft {i.e., low-carbon) steel (§ 314). 

308. What is steel? — We have just seen that cast iron is 
hard and brittle, whereas wrought iron is soft and tough. 
Also, that cast iron can be easily melted and poured into 

molds, whereas wrought 
iron softens readily and 
can be welded. More¬ 
over, cast iron contains 
a relatively high per 
cent of carbon (3 to 5), 
but wrought ifon a low 
per cent (0.15). Be¬ 
tween these extremes 
of composition come the 
different grades of steel. 

The physical proper¬ 
ties of the different 
grades of steel depend 
not only on the propor¬ 
tions of carbon, phos¬ 
phorus, silicon, sulfur, 
etc., but also to a large 
extent on the method 
of manufacture and 
treatment. 

309. Manufacture of steel. — Steel is made from cast iron. 
The aim in the manufacture is to prepare a product contain¬ 
ing the desired proportion of carbon but little or no sulfur, 
phosphorus, and silicon. The steel must also possess specific 
and known properties. This twofold aim is accomplished by 
several processes, often supplemented by special treatment. 



Fig. 132. — Sketch of a converter for mak¬ 
ing steel from cast iron by the Bessemer 
process. 
































IRON AND STEEL —IRON COMPOUNDS 239 


310. The Bessemer process. — This process consists in 
(1) burning out most of the impurities in cast iron by an air 
blast, and then (2) adding carbon and manganese to produce 
steel of the desired composition. 

It is carried on in a converter (Fig. 132). This is a huge, 
pear-shaped vessel, supported on trunnions so it can be tipped 
into different positions; one trunnion (A) is hollow, and at 
the bottom there are holes (B, B), through which a powerful 
blast of air can be blown. The converter is made of thick 
wrought iron plates and 
is lined with an infusible 
mixture, largely silica 
(sand, Si 02 ) with a little 
clay. 

When in use the con¬ 
verter is swung into a 
horizontal position, and 
fifteen to twenty tons of 
molten cast iron are run 
in, often directly from 
the blast furnace. The 
air blast is turned on, 
and the converter is 
swung back to a ver¬ 
tical position. As the 
air is forced in fine jets 
through the molten metal, the temperature rises, and the 
carbon, silicon, and manganese that are iii the iron are 
oxidized. The carbon forms carbon monoxide, which burns 
at the mouth of the converter in a large brilliant flame 
(Fig. 133), while the other oxides pass into the slag. This 
oxidation generates enough heat to keep the metal melted. 
In about fifteen minutes the diminished flame of burning 
carbon monoxide shows that the carbon has been oxidized 
and the other impurities removed. Then sufficient spiegel 
iron or ferro-manganese is added to the molten iron to 
furnish the proper amount of carbon (0.1 per cent for soft 
steel and 1.5 for hard, or 0.46 on the average) and manga¬ 
nese (about 0,9 per cent). By adding certain metals, e.gr., 



Fig, 133. — Making steel by the Bessemer 
process. Converter in operation. 





240 


A BRIEF COURSE IN CHEMISTRY 


aluminum, or iron alloys, e.g., ferro-titanium or ferro-silicon, 
inclosed gases are removed (by reacting with the “ deoxi¬ 
dizer ’’ or “ scavenger and a better quality of steel is 
produced. 

After the completion of the whole operation the converter 
is tilted and the metal is poured into ladles, and then into 
molds to form blocks called ingots (Fig. 134, left foreground), 
which are subsequently shaped into rails or other objects. 

The process described 
in the preceding para¬ 
graphs is called the acid 
Bessemer process be¬ 
cause the converter is 
lined with silica, which is 
an acid anhydride {i.e., 
the anhydride of silicic 
acid, H 2 Si 03 ). By(l)of 
this process the carbon 
and silicon can be re¬ 
moved but not all the 
sulfur and phosphorus. 
Both are objectionable. 
Sulfur makes steel brittle 
when hot, and phos¬ 
phorus, when cold. The 
acid Bessemer process is 
used in the United States 
because most domestic ores are low in phosphorus and 
sulfur. 

But in Europe the Thomas-Gilchrist or basic process is 
used. The converter in this process is lined with burned 
dolomite (^.6., practically a mixture of lime and magnesia, 
which are basic oxides), which removes the phosphorus and 
sulfur. This lining after use is known as Thomas slag. It 
is utilized as a fertilizer on account of its phosphorus (§ 300). 

311. The open-hearth process. — This process is used 
extensively in the United States and yields over 75 per cent 
of the annual production. It consists in heating a mixture 
of cast iron^ scrap iron, and iron oxide in a furnace lined with 








IRON AND STEEL —IRON COMPOUNDS 241 


a reacting material — burned dolomite or sand (see below). 
The process is conducted in a special kind of furnace called 



Fig. 135. — General view of a row of open-hearth furnaces. 



an open-hearth furnace (Fig. 135). A vertical section is 
shown in Fig. 136. The depressed hearth {H), on which the 
charge is put, is lined with burned dolomite in the basic 
process and with 
sand in the acid 
process. A slop¬ 
ing roof of fire 
brick is just 
above the hearth. 

At the base of the 
furnace are dupli¬ 
cate chambers of 
checkerwork (A 
and B, C and D) 
arranged for al¬ 
ternate use. 

To operate, 
fuel gas (or oil) 

is burned in a Fig. 136. — Sketch of an open-hearth furnace. 

furnace and the 

hot gases are passed through A, B, to the chimney, thus heat¬ 
ing this checkerwork very hot. The hot gases are then passed 










































































242 


A BRIEF COURSE IN CHEMISTRY 



through B and air through A, and the two brought together 
over the hearth. Here the gas burns and produces a high 

temperature on the 
hearth. Meanwhile the 
hot products of combus¬ 
tion and the unused gases, 
instead of escaping im¬ 
mediately up the chim¬ 
ney, are directed by 
valves through the other 
two chambers of the 
checkerwork C, D, and 
heat them. That is, 
while A and B are cool¬ 
ing, C and D are heating. 
At the proper time the 
fuel gas and air are 
shifted (by valves) to 
C, D and made to pass 
through this checkerwork to the hearth and out over the 
other checkerwork {A, B) to the chimney. 


Fig. 137. — Pouring molten iron into an 
open-hearth furnace. 



Fig. 138. — Putting an ingot of cast iron into an open-hearth furnace. 

By this plan the process is alternated, one checkerwork 
being cooled as the other is heated, and vice versa. It is only 







IRON AND STEEL —IRON COMPOUNDS 243 




Fig. 139. — Drawing a sample of steel from 
an open-hearth furnace for a chemical 
test in the laboratory. 


by this regenerative process, as it is called, that enough heat 
is obtained continuously to keep the charge melted as it 
becomes purer and purer. 

The charge consists of 
50 to 75 tons of cast iron 
and some iron oxide {e.g., 
hematite); scrap iron, 
or steel, and lime are 
usually added (Figs. 137, 

138) . The carbon is 
largely oxidized to carbon 
dioxide, which escapes. 

The silicon, sulfur, and 
phosphorus are con¬ 
verted into oxides (acid 
oxides), which form a 
slag with the hearth 
lining (a basic oxide). 

The charge is heated 
about eight hours. Samples are taken out at intervals (Fig. 

139) and tested by the chemist. When tests show that the 
metal contains the desired proportion of carbon and other 

constituents, the 
steel is tapped into 
ladles (Fig. 140), 
and certain mate¬ 
rials are added, e.g., 
aluminum or ferro¬ 
alloys of silicon, 
manganese, or tita¬ 
nium, to remove 
ob j ectionable gases. 
It is then quickly 
poured into molds 
and allowed to 
cool into ingots 
(Fig. 134). Subse¬ 
quently the ingots 
Fig. 140. — Tapping an open-hearth furnace. are softened by 












244 


A BRIEF COURSE IN CHEMISTRY 


reheating, and rolled, pressed, or stamped into desired 
shapes (Figs. 141, 142, 143). 

The open-hearth process is replacing the Bessemer process 
because it is easily controlled, operates on a large scale, can be 
followed by tests, and yields a more uniform product of the 

desired composition. 

Open-hearth steel is 
tough and elastic and 
is made into bridges, 
shafts, girders, heavy 
rails, large machines, 
large guns, and gun 
carriages. 

312. The crucible pro¬ 
cess. — This consists in 
melting wrought iron, or 
low carbon steel, in cruci¬ 
bles made of graphite 
and clay. The charge 
varies with the product 
desired. For example, 
iron oxide may be added, 
or certain metals, such 
as chromium, vanadium, 
or molybdenum. 

The crucibles are covered and then heated to a high temperature from 
four to five hours. During the heating the iron is slowly changed into 
steel by absorbing the proper proportions of carbon (0.75 to 1.5 per cent). 

Crucible steel is very hard, and is used to make tools, razor blades, 
files, knives, springs, drills, dies, pens, and needles. 

313. The electric process. — This process consists in heating a 
selected charge of cast iron or steel in an electric furnace (Fig. 144). 

The heat is generated, usually, by the arc established between the 
carbon electrode and the melted charge. Very high temperatures 
are obtained. Moreover, the operation is conducted in a non-oxidizing 
atmosphere. Hence, the electric process is advantageously used in 
conjunction with the Bessemer and open-hearth processes to produce a 
superior steel free from sulfur and phosphorus. 

314. Properties of steel. — The properties are numerous 
because there are many varieties of steel. Thus, steel, using 
this term broadly, is fusible and malleable, and can be 
forged, welded, and cast. Varieties containing 0.2 per cent 



Fig. 141. — Making a steel rail — first stage. 





IRON AND STEEL —IRON COMPOUNDS 245 


Fig. 142. — Making a steel rail — final stage. 


of carbon (“ low carbon are much like wrought iron and 
are called soft or mild steel. Structural steel contains more 
carbon (0.2 to 0.8 
per cent) and is hard 
like cast iron, while 
tool steel, which con¬ 
tains upwards of 1.5 
per cent of carbon 
(‘^high carbon’’), is 
very hard indeed. 

The properties of 
steel depend not only 
on the proportion of 
carbon (and other 
elements) but also 
on the special heat 
treatment which it 
receives. If steel 
containing the aver¬ 
age per cent of carbon (0.90) is heated to about 800° C. and 
then suddenly cooled by plunging it into cold water or oil, it 

becomes brittle and very 
hard (Fig. 145). But if 
the same kind of steel is 
heated and then cooled 
slowly, it becomes tough, 
elastic, and soft enough 
to be filed. All grades 
of hardness can be ob¬ 
tained between these 
two extremes. Thus, if 
hardened steel is re¬ 
heated to a definite tem¬ 
perature, determined ap¬ 
proximately by the color 
the oxidized metal as¬ 
sumes, and then properly 
cooled, almost any degree of hardness and elasticity can be 
obtained. This last operation is called tempering. 


Fig. 143. — Steel plate passing through 
the last rolling machine. 






246 


A BRIEF COURSE IN CHEMISTRY 


1:1 ni I I 


316. Special steels. — Special steels, or steel alloys as 
they are sometimes called, are made by melting a definite 

proportion of certain 


metals, or their ferro¬ 
alloys, into steel. These 
special steels have proper¬ 
ties which fit them for 
indispensable uses. 

1. Nickel steel con¬ 
taining up to 3.5 per cent 
nickel is hard, tough, and 
resists corrosion. It is 
used for armor plate, 
cables, drills, bridge 
trusses, propeller shafts, 
and marine engine parts. 
Steel containing 30 to 40 



Fig. 144. — Sketch of an electric furnace 
for the manufacture of steel. 


per cent of nickel resists corrosion so effectively it is made 
into pumps used for salt water. The alloy containing 36 per 
cent of nickel expands or contracts so very slightly it is 
made into surveyor’s tapes, pendulums, and scientific ap- 



Fig. 145. — Photomicrographs of hard steel (right) and cast steel (left). 
(Compare Fig. 131.) 

paratus. It is called invar. The alloy containing about 42 
per cent of nickel expands and contracts to the same ex¬ 
tent as platinum. It is used in making wire glass and as 


































IRON AND STEEL —IRON COMPOUNDS 247 


the leading-in wires of electric light bulbs instead of the 
scarce and expensive metal platinum. It is called platinite. 

2. Chromium steels, made by adding the alloy ferro- 
chrome to molten steel, contain from 1.5 to 2 per cent of 
chromium. They are very hard and are used for armor-pierc¬ 
ing projectiles, safes, and crushing machinery. If small 
proportions of other metals are added, such as nickel, vana¬ 
dium, and manganese, the steels are unusually tough, elastic, 
and hard. They resist shocks and strain, and are used in 
making automobile frames and axles. 

3. Tungsten steels are hard, and if of the right composi¬ 
tion, do not lose their hardness at a red heat. A typical one 
contains 18 per cent of tungsten, 3.5 per cent of chromium, 
and a small fraction of a per cent of vanadium. They are 
called high speed steel and are used to make tools for cutting 
metals. These tools, unlike the usual tool, do not lose their 
edge when heated but will cut when they are dull red from 
the heat produced by friction. Molybdenum sometimes 
replaces tungsten, and much less is needed to produce the 
same results as with tungsten. 

4. Manganese steels, containing 7 to 20 per cent of man¬ 
ganese, or about 13 per cent on the average, are extremely 
hard without being brittle. They are used for making the 
jaws of rock-crushing machinery, steam shovels, dredger 
buckets, safes, parts of brakes, and rails (especially railroad 
rails for curves). 

Stainless “ steel,” or stellite, is not steel at all, but an alloy of chro¬ 
mium (15 to 35 per cent), cobalt (50 to 75), and small per cents of iron, 
carbon, and tungsten. It does not rust, or even stain, and is used to 
make cutlery and surgical instruments. 

316. Rusting. — Iron and steel rust readily in moist air, 
i.e., they gradually change into a brittle, porous, loosely ad¬ 
herent brown solid. This solid, which is called iron rust, 
varies in composition, though it is often considered as essen¬ 
tially hydrated ferric oxide (3 Fe 203 • H 2 O or 2 Fe 203 .3 H 2 O). 

Rusting is a complex process and is explained in different 
ways. An acceptable interpretation is based on the ionic 
theory. The iron goes into solution as ferrous ions (Fe++). 


248 


A BRIEF COURSE IN CHEMISTRY 


Hydrogen ions (H+) from the water lose their charges and 
become atoms; the atoms unite to form molecules which 
escape as hydrogen gas (H 2 ). The ferrous ions combine with 
the hydroxyl ions (OH~) left in the water, thereby forming 
ferrous hydroxide (Fe(OH) 2 ), which is subsequently con¬ 
verted into the complex substance called iron rust. Once 
begun, rusting proceeds rapidly, because the film of rust is 
not compact enough to protect the metal. 

The loss from rusting is enormous. Various methods have 
been devised to prevent it. One is to expose the hot metal 



Courtesy Central Alloy Steel Co. 
Fig. 146. — Objects made of rustless iron. 


to steam until a thin film of magnetic oxide (Fe 304 ) forms 
on the surface. This film adheres firmly and prevents rusting 
to a large extent. Iron so coated is often called Russia iron. 

Another method is to paint the metal first with red lead 
paint and then with some less conspicuous kind. Sometimes, 
special black paints or aluminum (powder) paint are used, 
the latter having special value as a coating for oil storage 
tanks. 

Sometimes an enamel coating is melted upon the metal. 
This process is used in making kitchen utensils. 

Iron is often “ tinned,” i.e., sheets of iron or “ low. carbon ” 
steel are carefully cleaned and dipped into melted tin. 






IRON AND STEEL —IRON COMPOUNDS 249 



Tinned iron, or tin plate as it is often called, does not rust 
until the tin is worn off, and then the rusting proceeds rapidly. 

Another metallic coating largely used is zinc. The iron 
is cleaned and dipped into melted zinc. The product is 
called galvanized iron. It is extensively used for wire, net¬ 
ting, roofs, pipes, cornices, and water tanks. 

Special alloys of iron called rustless iron are a recent prod¬ 
uct (Fig. 146). One contains about 14 per cent of silicon. 


Courtesy Duriron Co. Inc. 

Fig. 146a. — Duriron waste pipe on the ceiling of a chemical laboratory for 
draining the sinks on the floor above. 

It is called duriron. It is used in acid plants and as the 
waste pipe in laboratories (Fig. 146a). 

317. Two series of iron compounds. — Iron forms two 
series of compounds — ferrous and ferric. Corresponding 
members differ in properties. In each series the iron plays 
a different r61e. The valence of iron is + 2 in ferrous com¬ 
pounds and -f- 3 in ferric. Ferrous salts in solution give 
ferrous ions (Fe++) and ferric salts give ferric ions (Fe+++). 

Ferrous compounds are readily oxidized into the corre¬ 
sponding ferric compounds. For example, ferrous chloride in 
hydrochloric acid solution becomes ferric chloride in the pres¬ 
ence of an oxidizing agent, such as nitric acid or potassium 







250 


A BRIEF COURSE IN CHEMISTRY 


permanganate. Whereas ferric compounds are readily re¬ 
duced to ferrous compounds. Thus, ferric chloride becomes 
ferrous chloride by the action of hydrogen sulfide, sulfurous 
acid, or stannous chloride (SnCh). So also ferrous chloride 
becomes ferric chloride by action with chlorine, while ferric 
chloride becomes ferrous chloride by action with iron, thus: — 

2 FeCb -I- CI 2 = 2 FeClg, and 2 FpCb + Fe = 3 FeCb 

That is to say, the change from a ferrous to a ferric com¬ 
pound is the same type of change as from sulfurous acid to 
sulfuric acid. It is an example of oxidation, while the reverse 
is reduction. Oxygen need not be involved. We often use 
the terms oxidation and rediLction in a broader sense than mere 
addition and removal of the element oxygen. 

Let us interpret the broader meaning of these terms. Oxy¬ 
gen belongs to the class of negative elements. So does 
chlorine. Now if we add chlorine to ferrous chloride, we 
are doing chemically just what we do if we add oxygen to 
sulfurous acid, viz., oxidizing. In a broad sense, then, oxida¬ 
tion is the process of adding oxygen, chlorine, or another 
negative element. Conversely, reduction is the process of 
removing oxygen, chlorine, or another negative element. 

As stated above, the valence of iron is + 2 in ferrous and 
-f 3 in ferric compounds. In passing from ferrous to ferric 
compounds the valence of the iron increases; and conversely, 
from ferric to ferrous it decreases. From the standpoint of 
valence, oxidation is an increase in the valence of an element, 
whereas reduction is a decrease. 

Furthermore, since dissolved iron salts pass readily from 
one series to the corresponding members of the other, oxida¬ 
tion is sometimes called the gain of a positive charge, e.g., 
Fe++ to Fe"^^, whereas reduction is the loss of a positive 
charge, e.g., Fe+++ to Fe++. 

Still another way to describe oxidation and reduction is in 
terms of electrons (§§ 357, 358, 424, 529). Oxidation is the 
loss and reduction the gain of electrons, thus: — 

3 Fe++- >■ 2 Fe+++ and 2 Fe+++- >■ 3 Fe++ 

Oxidation (loss of E) Reduction (gain of E) 


IRON AND STEEL —IRON COMPOUNDS 251 


318. Iron hydroxides. — The white solid formed by the 
interaction of a ferrous salt and an hydroxide ^is ferrous 
hydroxide (Fe(OH) 2 ). Exposed to the air, it soon turns 
peen, and finally brown, owing to oxidation to ferric hydrox¬ 
ide. Ferric hydroxide (Fe(OH) 3 ) is a reddish brown solid, 
formed by the interaction of a ferric salt and an hydroxide, 
thus: — 

FeCls + 3NaOH = Fe(OH )3 + 3 NaCl 

Ferric Sodivtm Ferric Sodium 

Chloride Hydroxide Hydroxide Chloride 

319. Ferrous sulfate. — This is a green salt obtained by 
the interaction of iron (or of ferrous sulfide) and dilute sul¬ 
furic acid. It is manufactured by roasting {i.e., oxidizing) 
iron pyrites (FeS 2 ), or by exposing pyrites to moist air; the 
mass is extracted with water containing scrap iron and a little 
sulfuric acid. Considerable is obtained in “ pickling iron 
castings with dilute sulfuric acid. The large light green crys¬ 
tals (FeS 04 .7 H 2 O) are called green vitriol or copperas. 

Large quantities are used in dyeing silk and wool, as a 
disinfectant, and in manufacturing ink, bluing, and pigments. 
Much black writing ink is made essentially by mixing ferrous 
sulfate, nutgalls, gum, and water. A mixture of ferrous 
sulfate and lime is sometimes used to purify water and sewage 
by the settling process (§81). 

320. Ferrocyanides and ferricyanides. — The most im¬ 
portant are potassium ferro- and ferricyanides. Potassium 
ferrocyanide (K 4 Fe(CN) 6 ) is a lemon-yellow solid. The 
crystallized salt contains three molecules of water of hydra¬ 
tion. It is sometimes called yellow prussiate of potash. 
Unlike the simple cyanogen compounds (e.gr., HCN and 
KCN), it is not poisonous. Potassium ferricyanide 
(K 3 Fe(CN) 6 ) is a dark red crystallized solid (without water 
of hydration). It is often called red prussiate of potash. 

321. Making a blue print. — Blue print paper is prepared 
by coating paper with a mixture of potassium ferricyanide 
and ammonium ferric citrate solutions, and drying in a dark 
place. In the sunlight the ferric salt is partly reduced and 
forms a bronze-colored deposit by interaction with the potas¬ 
sium ferricyanide. 



252 


A BRIEF COURSE IN CHEMISTRY 


If such prepared paper is covered with a photographic 
negative, or with transparent cloth marked with lines in 
black ink, and exposed to the sun or an electric light, the paper 
is acted upon only in the exposed places. Upon washing, 
the exposed parts become blue, and the covered parts white 
— hence the name blue print.” 

322. Tests for ferrous and ferric compounds. — Solutions 
of ferrous salts and potassium ferricyanide interact and pre¬ 
cipitate ferrous ferricyanide (Fe 3 (Fe(CN) 6 ) 2 ). This is a 
dark blue solid, called TurnbulFs blue. Ferric salts interact 
with potassium ferrocyanide and precipitate ferric ferrocyan- 
ide (Fe 4 (Fe(CN) 6 ) 3 )- This precipitate is likewise dark blue 
and is called Prussian blue. By these tests ferrous and ferric 
salts can be distinguished. Hence to test for — 

Ferrous salts add potassium ferricyanide. 

Ferric salts add potassium ferrocyanide. 

In each test we obtain a dark blue precipitate, but only if 
we use “ ferrous with ferri- ” and ferric with ferro-.” 

Besides the above tests for ferric salts, potassium thio¬ 
cyanate (KCNS) produces a red solution of ferric thiocyan¬ 
ate (Fe(CNS) 3 ) with ferric salts, but leaves ferrous salts un¬ 
changed. 

Since these two blue ” tests are apt to confuse, let us 
interpret them. Ferrous salts in solution yield ferrous ions 
(Fe++). Potassium ferricyanide in solution yields potassium 
ions (K+) and ferricyanide ions (Fe(CN)6) . When fer¬ 

rous chloride and potassium ferricyanide solutions are 
mixed, ferrous ions unite with ferricyanide ions and form 
ferrous ferricyanide, thus : — 

3 Fe++ + 6 Cl- + 6 K+ + 2 Fe(CN)6- 

= Fe3(Fe(CN)6)2 + 6 Cl" + 6 K+ 

Ferrous Ferrricyanide 

Similarly, ferric chloride and potassium ferrocyanide react 
thus:— 

4 Fe+++ + 12 Cl- + 12 K+ + 3 Fe(CN)6- 

= Fe4(Fe(CN)6)3 + 12 01“ + 12 K+ 

Ferric Ferrocyanide 


IRON AND STEEL —IRON COMPOUNDS 253 


The dark blue precipitate is formed only by the combination 
of either (1) ferrous and ferricyanide ions or (2) ferric and 
ferrocyanide ions. 


EXERCISES 

1. Describe the manufacture of cast iron. 

2. Apply Exercise 1 to wrought iron. 

3. State the physical properties of (a) cast iron and (6) wrought iron. 

4. Describe the manufacture of steel by (a) the Bessemer process, 
(6) the open-hearth process, (c) the crucible process, and (d) the electric 
process. 

5. How are ferrous compounds changed into ferric compounds, and 
vice versa ? Give equations. 

6 . Write the formulas of (a) the ferrous and (6) the ferric salts of the 
following acids: carbonic, nitric, sulfuric, orthophosphoric (H3PO4). 

7. Complete and balance the following: (a) FeCL + (NH 4 ) 2 S 

= FeS +-; (b) FeCh +- = Fe(OH )3 + NH4CI; (c) Fe + O 

= FeaOs; (d) K 3 Fe(CN )6 +- =-+ K 2 SO 4 ; (e) K 4 Fe(CN )6 

+- =-+ KCl. 


PROBLEMS 

(See Problems at the end of Chapters VIII, IX.) 

1. Calculate the weight of iron in 70 tons of hematite (95 per cent 
pure). 

2. What is the simplest formula of a compound, if 9 gm. of it yielded 
4.8 gm. of sulfur and the rest iron? 


SUGGESTIONS FOR LABORATORY WORK 

(References are to Newell’s Laboratory Exercises in Chemistry) 

Exercise 59 — Tests for Iron Salts. 

Exercise 60 — Reduction and Oxidation of Iron Salts. 
Exercise S5 (a) — Slow Oxidation — T. 

Exercise *2 — Mixture and Compound. 


SUPPLEMENTARY SECTIONS FROM PART II 

At this point sections may be selected from Topics XXII (Copper), 
XXIII (Magnesium — Zinc — Mercury), XXIV (Silver — Photog¬ 
raphy — Gold), XXV (Aluminum — Clay Products^ XXVI (Lead), 
XXVII (Radium — Radioactivity). 






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BRIEF COURSE IN CHEMISTRY 


PART II 

SUPPLEMENTARY TOPICS 

SUITABLE FOR SPECIAL NEEDS AND 
FOR ADDITIONAL STUDY 










SUPPLEMENTARY TOPICS 

TOPIC I: ELEMENTS —ENERGY —CHEMICAL CHANGE 

326. Distribution of the elements. — Approximately 98 
per cent of the outer shell (about ten miles deep) of the earth’s 
crust consists of compounds of only eight elements. The 
per cent of these elements is shown in Table VIII. 


TABLE VIII. — Per Cent op Combined Elements 
IN THE Earth’s Crust 


Oxygen . . . 46.71 

Silicon . . . 27.69 

Aluminum . . 8.07 


Iron .... 5.05 
Calcium . . . 3.65 

Sodium . . . 2.75 


Potassium . . 2.58 

Magnesium . . 2.08 

Remainder . . 1.42 


Compounds in the ocean contain only a few elements, as 
seen by Table IX. 


TABLE IX. — Per Cent of Combined Elements in 
THE Ocean 


Oxygen . . . 85.79 

Hydrogen . . 10.67 

Chlorine . . 2.07 


Sodium . . . 1.14 

Magnesium. . 0.14 

Sulfur . . . 0.09 


Calcium . . 0.05 

Bromine . . 0.008 

Carbon . . . 0.002 


The atmosphere contains about 21 per cent of oxygen, 78 
of nitrogen, and 1 of argon — all as free elements. 

The human body, complicated as it is, consists mainly of 
compounds of about six elements together with as many 
more elements in very small proportions, some only in traces 
(Table X). 


257 















258 


A BRIEF COURSE IN CHEMISTRY 


TABLE X. — Per Cent (Average) of Combined 
Elements in the Human Body 


Oxygen . . 

. 65.00 

Phosphorus. 

. 1.00 

Magnesium 

. 0.05 

Carbon . . 

. 18.00 

Potassium . 

. 0.35 

Iron . . . 

. 0.004 

Hydrogen . 

. 10.00 

Sulfur . . 

. 0.25 

Iodine . . 

. trace 

Nitrogen 

. 3.00 

Sodium . . 

. 0.15 

Fluorine 

. trace 

Calcium 

. 2.00 

Chlorine. . 

. 0.15 

Silicon . . 

. trace 


326. A good example of a chemical change. — If we light 
a candle and let it burn, the candle wax slowly disappears — 
“ burns away/’ By holding a cold, dry bottle over the burn¬ 
ing candle (Fig. 147), a film of water 
gathers on the inside. Now if we re¬ 
move the bottle, pour in some calcium 
hydroxide (limewater), and shake the 
bottle, the clear liquid becomes cloudy, 
owing to the formation of fine parti¬ 
cles of calcium carbonate. These were 
formed by a chemical change between 
the added calcium hydroxide and the 
carbon dioxide (produced by the burn¬ 
ing candle). Hence, the candle wax 
did not really burn away.” By 
burning in the air, the wax changed 
into water and carbon dioxide. 

327. Energy and chemical change. — One striking feature 
of the experiment in which iron and sulfur are heated can 
not escape observation. That is, the reaction, once started 
by heat, proceeds with the evolution of a great deal of heat. 
In this experiment part of the chemical energy in the iron and 
the sulfur was transformed into heat. All chemical reactions, 
of course, do not result in the conspicuous evolution of heat, 
indeed in many reactions heat is absorbed, but every chemi¬ 
cal change involves some change in energy. So we say: — 

A characteristic of chemical change is the transformation of 
chemical energy into heat energy, or light energy, or both, 
and sometimes into electrical energy. 



Fig. 147. — Studying the 
chemical change in a 
burning candle. 




















THE METRIC SYSTEM 


259 


TOPIC II: THE METRIC SYSTEM 

328. The metric system. — The fundamental unit of this 
system is the meter. It is the unit of length and is 39.37 
inches. The meter and the other units have multiples and 
submultiples, which are designated by prefixes attached to 
the particular unit. . Thus, kilo- is equivalent to 1000, and 
deci-, centi-, and milli- to 0.1, 0.01, and 0.001. Hence : — 

10 millimeters (mm.) = 1 centimeter (cm.) 

10 centimeters = 1 decimeter (dm.) 

10 decimeters = 1 meter (m.) 

In chemistry, the measures of length usually used are the 
millimeter (mm.) and centimeter (cm.); occasionally they 


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INCHES 

WTT 

4 

CENTIMETERS 

lllllllllllllllllll llllllllllllllllllilllllllllllllllllllllllllllllll^^^ 


Fig. 148. — Comparative scale — metric below, English above. On the 
metric scale the numbered divisions are centimeters and the smallest are 
millimeters. 


meter (m.) is used. A comparative scale (metric and 
English) is given in Fig. 148. Note that 2.5 cm. = 1 in. 
(very nearly). For convenience we might remember that 
a five cent piece is 2 centimeters in diameter and a worn ten 
cent piece is 1 millimeter thick. 

The unit of weight is the gram. It is the one-thousandth 
part of a standard weight called the kilogram. Thus: — 

10 milligrams (mg. or mgm.) = 1 centigram (eg. or cgm.) 

10 centigrams (eg. or cgm.) = 1 decigram (dg. or dgm.) 

1000 grams (gm.) = 1 kilogram (kg. or kgm.) 

A kilogram weighs about 2.2 pounds (avoir.). Large 
weights are often expressed in kilograms (and a decimal 
fraction). A gram weighs 15.43 grains, or also nearly one 
thirtieth of one ounce (1 ounce = 28.55 grams). Small 
weights are expressed in grams and a decimal fraction. Thus, 




260 


A BRIEF COURSE IN CHEMISTRY 


2 grams, 2 centigrams, and 5 milligrams is written 2.025 
grams (or gm.). A pint of water weighs a little less than 500 
grams. A five cent piece weighs 5 grams. 

The unit of volume is the liter. It is equal to the capacity 
of the vessel containing a kilogram of water. A liter equals 
about one quart. The table is: — 

1000 cubic millimeters = 1 cubic centimeter (cc). 

1000 cubic centimeters = 1 cubic decimeter (cu. dm.) 

1 cubic decimeter = 1 liter (1.) 

In chemistry the cubic centimeter and the liter are the de¬ 
nominations used to measure and express volume. Thus, a 
test tube measuring 15 X 1.8 cm. holds about 30 cc., a large 
test tube (20 X 2.5 cm.) holds about 75 cc., and the large 
greenish glass acid bottles hold 2.5 1. 

An approximate relation which can be used for water 
(and other liquids of the same specific gravity) is 1 liter = 1 
kilogram = 1 cubic decimeter = 1000 cubic centimeters = 
1000 grams = 2.2 pounds (avoirdupois). 


TOPIC III: OXYGEN 

329. Preparation of oxygen from various substances. — 

Oxygen can be prepared from other compounds besides 
potassium chlorate. 

(1) Mercuric oxide (HgO), barium dioxide (Ba02), lead 
dioxide (Pb 02 ), and lead tetroxide (Pb 304 ), if heated sepa¬ 
rately, liberate oxygen and may be used to prepare the gas 
in the laboratory. 

(2) When water is dropped slowly upon sodium peroxide 
(Na202), oxygen is liberated. The experiment must be done 
on a small scale. The equation is: — 

2Na202 + 2 H 2 O = O 2 + 4NaOH 

Sodium Peroxide Water Oxygen Sodium Hydroxide 

(3) A mixture of hydrogen peroxide (H 2 O 2 ), sulfuric acid, 
and potassium permanganate (KMn 04 ) solution yields oxy¬ 
gen. 


OXYGEN 


261 


330. Lavoisier’s famous experiment on combustion. — 

Lavoisier showed that oxygen is necessary for combustion. 

He put mercury in a retort having a long neck which com¬ 
municated with a jar (Fig. 149). The jar contained air con¬ 
fined over mercury (in the lower vessel). He heated the re¬ 
tort and kept the mercury in the retort just below the boiling 
point for twelve days. When the apparatus was cold, he 
noticed two things: (1) A red substance had accumulated 
on the surface of the mercury in the retort, and (2) the air 
in the retort and jar had decreased in volume about one fifth. 

He picked off the red substance, heated it intensely, col¬ 
lected the liberated gas, and measured its volume. It was 
the same as the decrease in 
the volume of the original air 
inclosed in the apparatus. He 
tested this gas, and found it 
had the properties of the gas 
previously obtained by 
Priestley, especially the prop¬ 
erty of making a lighted 
candle burn vigorously. He 
also tested the gas left behind 
in the retort and jar and noted 
that it extinguished a lighted 
candle. From these facts Lavoisier concluded: (1) air con¬ 
tains a gas which is removed by heating substances in it, 
(2) this gas is about one fifth of air, (3) this gas unites with 
substances in the process called combustion or burning, and 
(4) this gas is identical with the gas (f.e., oxygen) obtained 
by Priestley from his red precipitate of mercury.^’ 

331. Oxidation and energy. — During oxidation, some of 
the chemical energy of the oxygen and the other substances 
is transformed into heat energy and often into light energy. 
Oxidation takes place at different rates. But the actual 
amount of heat energy liberated in a single case of oxidation 
is the same, whether the oxidation is slow or rapid (§ 333). 
To be sure the temperature resulting from slow oxidation is 
much lower than from rapid oxidation. But temperature, 
as registered on a thermometer, tells how hot ’’ a substance 



Fig. 149. — Apparatus used by La¬ 
voisier in his famous experiment 
on combustion. 





















262 


A BRIEF COURSE IN CHEMISTRY 


is. Quantity, ^.e., ‘‘ how much ” heat, is measured in a calo¬ 
rimeter (§ 384) and is stated in calories. A small calorie 
(cal.) is the amount of heat required to raise the temperature 
of one gram of water one degree centigrade (usually from 15° 
to 16° C.). 


TOPIC IV: CARBON 

332. How coal was formed. — Ages ago the vegetation 
was exceedingly dense and luxuriant upon the land slightly 



Fig, 150. — Section of the earth’s crust showing layers of coal. 


raised above the sea. In process of time this vegetation 
decayed, accumulated, and became covered with sand, mud, 
and water. The vegetable matter was slowly changed by 
heat and pressure into impure carbon. These geological 
and chemical changes were repeated, and as a result there 



Fig. 151. — Fossil found in a coal Fig, 152. — Section of coal as seen 
bed. through a microscope. 


are in the earth layers or seams of carbonaceous matter vary¬ 
ing in thickness and composition (Fig. 150). These are the 




CARBON 


263 


coal beds. Coal beds contain proofs of their vegetable origin, 
viz., impressions of vines, stems, and leaves of plants, and 
similar vegetable substances (Fig. 151). A thin section of 
coal examined through a microscope shows a vegetable struc¬ 
ture (Fig. 152). 

333. Carbon and energy. — Careful experiments made in a 
calorimeter (§ 384) show that the same amount of heat, stated 
in calories (§§ 235, 383), is produced in 
a given oxidation reaction, provided, of 
course, the same weight and kind of 
product is formed in each case by oxi¬ 
dizing a certain weight of substance. 

For example, 12 gm. of carbon (char¬ 
coal) form 44 gm. of carbon dioxide. 

Now this weight of carbon (12 gm.) 
in burning liberates 97,000 calories, 
whether we burn it slowly or rapidly. 

The thermo-chemical equation for this 
reaction is: — 

C + 02= CO 2 + 97,000 calories 

12 gm. 32 gm. 44 gm. 

This equation may be interpreted as fol¬ 
lows: When carbon and oxygen react, 
carbon is oxidized; some of the chem¬ 
ical energy in these substances is trans¬ 
formed into heat energy; and 97,000 
calories are liberated for every 12-gram 
portion of carbon which is oxidized. 

334. Carbon dioxide and plants. — The fact that plants 
take up carbon dioxide and reject oxygen can be readily illus¬ 
trated. 

Fresh green leaves are put into the flask (Fig. 153), which is then 
completely filled with water saturated with carbon dioxide. The 
stopper with its funnel is pushed in to exclude the air, the funnel is partly 
filled with the same liquid, and the test tube is filled and arranged as 
shown in the figure. On exposure to the sunlight for several hours, a 
gas collects in the test tube. The usual test shows that the gas is oxygen. 



Fig. 153.—Experiment 
showing the absorp¬ 
tion of carbon dioxide 
and liberation of oxy¬ 
gen by plants. 
















264 


A BRIEF COURSE IN CHEMISTRY 


TOPIC V: HYDROGEN 

335. Hydrogen and energy. — The burning of hydrogen 
in oxygen involves the transformation of a large amount of 
chemical energy into heat energy. Thus, 1 gm. of hydrogen 
in burning to water (vapor) yields 29,050 calories; if the re¬ 
sulting water vapor is condensed to the liquid form, then 
about 34,000 calories are liberated. This is a large amount of 
heat energy. One gram of carbon, for example, gives only 
about 8000 calories (§ 333). The thermo-chemical equation 
may be written in full thus: — 

2 H 2 + O 2 = 2 H 2 O + 116,200 calories 

336. Hydrogenation of fats. — Liquid fats (usually called 
oils) are mixtures of several fats. One of the ingredients is a 
liquid called olein. It contains 2 hydrogen atoms less in a 
molecule than the solid fat stearin. The chemical change 
hastened by the catalyst is the chemical addition of 2 hydro¬ 
gen atoms to each molecule of olein. 


TOPIC VI: GASES AND THEIR MEASUREMENT 

337. Gases. — Gases are sensitive substances. They 
contract and expand with the slightest change in temperature 
or pressure. Hence the volume of a gas changes readily. 

338. Measuring gases. — In chemistry we often need to 
know the weight of a certain volume of a gas, e.g., 1 liter, or 
22.4 liters, of a gas, or to compare the weights of the same 
volume of different gases. But since gases change so readily 
in volume, we can not compare weights of different gases 
unless the volumes are measured at the same temperature 
and pressure. 

We can find the volume of a gas by measuring the capacity 
of the container; or we can conduct the gas into vessels on 
which the capacity is marked in some conspicuous way 
(Fig. 154). We can measure volumes very accurately. And 
if we find (or know) the volume of a sample of a gas, we can 
find its weight by calculation (§ 345). 


GASES AND THEIR MEASUREMENT 


265 


339. Finding the volume of a gas. — As already stated, 
the volume of a gas changes readily with variations of temper¬ 
ature and pressure. This means that the actual volume of a 
definite weight of gas depends on the temperature and pres¬ 
sure prevailing at the time the volume is read. Hence, to 
give an accurate meaning and value to the measurement of a 
gas volume we must specify the exact tempera¬ 
ture and pressure at which the measurement was 
made. 

By agreement among scientists, the normal or 
standard temperature is 0° C., ^.e., zero degrees 
on the centigrade thermometer, and the normal 
or standard pressure is 760 mm., i.e., the pres¬ 
sure when the mercury column of a barometer 
is 760 millimeters high. These conditions of 
temperature and pressure (0° C. and 760 mm.) 
are called standard conditions. Under these 
conditions a liter of dry oxygen gas, for exam¬ 
ple, weighs 1.429 gm. But at another temper¬ 
ature or pressure the weight of oxygen gas filling 
a liter vessel would be different. Thus, if the 
pressure is increased, the volume becomes less, 
more gas must be added to bring the volume 
up to a liter, and this total quantity of oxygen 
would weigh more than 1.429 gm. That is, a 
liter vessel, when full, contains a liter, but the weight of the 
contents varies with the quantity of gas in the vessel. 

If we could measure gases at 0° C. and 760 mm., their 
volumes could be compared directly, and the weights deduced 
or obtained directly from these volumes would be a true 
measure of the actual quantity of the gases in the observed 
volumes. But it is not convenient to measure gases at 0° C. 
and 760 mm. So it is customary to measure the volume at 
the temperature and pressure prevailing at the time of the 
experiment, and then compute what the observed volume would 
occupy if it were under standard conditions. This mathemati¬ 
cal process is possible because two laws define the behavior of 
gases under variation of temperature and pressure. These 
laws are the law of Charles and the law of Boyle. 


Fig 


154 . — 
Graduated 
tube for 
measuring 
the volume 
of a gas. 




266 


A BRIEF COURSE IN CHEMISTRY 


CENTIGRADE ABSOLUTE 
Water 


100 ® 


Boils 


Freezes 


340. The relation between volume and temperature, and 

the law of Charles. — Gases expand when heated and con¬ 
tract when cooled. The change in volume is the same for 
all gases, provided there is no change in pressure. Thus, at 
constant pressure a gas at 0° C., if warmed to 1° C., will 
expand of the volume it occupied at 0° C. And for each 
additional rise of 1° C., the volume will 
increase of the volume at 0° C. 
Conversely, if a gas at 0° C. is cooled to 
— 1°C., it will contract of its vol¬ 
ume, and for each additional fall of 
1° C. there will be a decrease of of 
the volume at 0° C. 

At this rate of contraction a gas 
would have no volume at — 273° C. I 
But this rather startling result does 
not actually happen, because all gases 
become solids before the temperature 
reaches — 273° C. Nevertheless for 
convenience we use this point (— 273°) 
on the centigrade thermometer as the 
starting point of another temperature 
scale called the absolute temperature. 
The temperature — 273° C. is called 
absolute zero, and the temperatures 
reckoned from this point as zero are 
called absolute temperatures. If an 
“ absolute thermometer ” were con¬ 
structed with divisions equal to those 
on the centigrade scale and compared 
with a centigrade thermometer, the 
scales would be seen to be simply related (Fig. 155). If we' 
designate absolute temperatures by A. and centigrade tem¬ 
peratures by C., then 373° A. is 100° C., and 253° A. is 
— 20° C. That is, absolute degrees are obtained by adding 
273 algebraically to centigrade degrees. 

How is absolute temperature used ? Suppose we have 273 
cc. of gas at 0° C. This volume becomes 274 cc. at 1° C. 
and 280 at 7° G. Now 0° C., 1° C., and 7° C. are 273° A., 


-182.5® 


Boils 


90.5® 


-273® 


Absolute 


0 ® 


Zero 

Fig. 155. — Centigrade 
and absolute scales. 





GASES AND THEIR MEASUREMENT 


267 


274 A., and 280° A., respectively. That is, stated in general 
terms as the law of Charles, we can say: — 

A.t constdnt pvessuve the volumes of cl given sample of gas 
at different temperatures vary directly as the absolute tempera¬ 
tures. 

341. How we apply the law of Charles. — If we have 100 
cc. of oxygen at 25° C. and wish to know what the volume 
would be at 0° C., we do not have to cool the gas to 0° C. 
and then measure the volume. We know; (1) that a gas 
contracts uniformly for each degree of fall in temperature, 
and (2) that the volumes vary directly as the absolute temper¬ 
atures. Hence we can compute the volume to which the 100 
cc. would contract if it were cooled from 25° to 0° C. The 
steps in the computation are these:— 

(а) Change each centigrade temperature to absolute tem¬ 
perature by adding 273. 

(б) Make a proportion involving the relation stated in 
(2), i.e., observed volume: required volume :: observed abso¬ 
lute temperature : required absolute temperature. 

(c) Solve the proportion for x. 

Thus: (a) 25 + 273 = 298, and 0 + 273 = 273; (6) 100 : 
X :: 298 : 273; (c) x = 100 X = 91.6. That is, 91.6 cc. 
is the volume 100 cc. would become if cooled from 25° C. to 
0° C. 

This form of computation is sometimes confusing. So a simpler one 
may be used, if desired. Since the temperature is lower, the volume 
will be smaller, because volumes change in the same way as the tempera¬ 
tures. Obviously, in the example given, 100 must be multiplied by a 
fraction having the numerator smaller than the denominator, i.e., 

100 X IH = 91-6. 

The mathematical process described in the above para¬ 
graphs is called correcting for temperature. 

342. The relation between volume and pressure, and the 
law of Boyle. —Many gases are collected in bottles or tubes 
filled with water. In measuring the volume of the gas, the 
levels of the water are made the same inside and outside the 
collecting vessel by raising or lowering this vessel. When the 
levels are the same, the gas is under atmospheric pressure, 
i.e., the pressure of the atmosphere exerted on the exposed 


268 


A BRIEF COURSE IN CHEMISTRY 


surface of the liquid is transmitted through the water to the 
gas. Hence the pressure which the gas is under when the 
volume is read is the pressure of the atmosphere. 

The pressure of the atmosphere is found by reading the 
barometer. A common form of barometer is shown in Fig. 

156. Near the top is a scale on which we can 
read the height of the mercury column. That 
is, by reading on the scale the height of the mer¬ 
cury — reading the barometer ” — we are 
measuring the pressure of the atmosphere at 
the time of the observation. The normal height 
of the barometer is 760 millimeters (mm.). The 
pressure of the atmosphere when the barometer 
is 760 mm. is called normal or standard pressure. 
It is also called one atmosphere. 

The effect of pressure on the volume of a gas 
was first studied by the English chemist Boyle. 
He found that the volume is halved when the 
pressure is doubled, and so on. That is, the 
volume and pressure are in an inverse relation. 
The general fact may be stated as the law of 
Boyle, thus: — 

The volume of a gas at constant temperature 
changes inversely with changes in pressure. 

343. How we apply the law of Boyle. — In 

most experiments in chemistry we actually read 
a gas volume at the prevailing pressure, but we 
usually want to know what the volume would be 
if the pressure were 760 mm. Since we know 
that volumes are inversely related to their 
corresponding pressures (§ 342), we can com- 
Fig. 156 .—a pute the volume a gas would have, if the pres- 
barometer. 'wqyq changed to 760 mm. The steps in 

the computation are these: — 

(а) Make a proportion involving the inverse relations 
of volumes and pressures, i.e., observed volume : required 
volume :: required pressure : observed pressure. 

(б) Solve the proportion for x. 







GASES AND THEIR MEASUREMENT 


269 


For example, suppose we wish to know what volume 100 cc. 
at 775 mm. would have at 760 mm. (a) 100 : x :: 760 : 775 ; 
(h) X = 100 X Yi% = 102. This result means that 102 cc. 
is the volume 100 cc. would become if the pressure were 
changed from 775 mm. to 760 mm. 

If the above form of computation is confusing, a simpler one may 
be used. Since pressure and volume are inversely related, the less the 
pressure the larger the volume. So the observed volume must be multi¬ 
plied by a fraction having the numerator larger than the denominator, 
i.e., 100 X^U = 102 cc. 

The mathematical process described in the above para¬ 
graphs is called correcting for pressure. 

344. How we correct a gas volume for both temperature 
and pressure. — Measuring gases usually involves both tem¬ 
perature and pressure. Hence we compute what the volume 
would be at 0° C. and 760 mm. by combining the two mathe¬ 
matical processes into a single calculation. This combined 
mathematical process is called reducing to standard condi¬ 
tions, and the final volume is called the corrected volume. 
Suppose the observed volume is 100 cc., the temperature 
25° C., and the pressure 775 mm. Then the volume at 
standard conditions is computed thus : — 

100 X m X = 93.4 


Sometimes the reduction to standard conditions is more conveniently 
performed by substituting the observed values in the corresponding 
mathematical formula: — 


T2 = 


Vi X Pi 

760(1 H-(0.00366 X t)) 


In this formula, which merely involves the mathematical processes just 
described, V 2 = final volume, Vi = observed volume. Pi = observed 
pressure, t = observed temperature (centigrade), and 0.00366 = 


345. Finding the weight of a given volume of a gas. — 

The weight of one liter of oxygen is 1.429 gm. The weight 
of one liter of oxygen is readily found by experiment. 

(a) The apparatus is shown in Fig. 157. A mixture of potassium 
chlorate and manganese dioxide is put in the test tube A. The part 
AF is weighed. The bottle B is filled with water. The empty bottle 



270 


A BRIEF COURSE IN CHEMISTRY 


D is weighed. The mixture is heated and oxygen forces water from B 
into D. When sufficient gas (about half the volume of B) has been 

liberated, the heating 


is stopped and the 
apparatus is allowed 
to cool. (6) The tem¬ 
perature and pressure 
are read, say 20° C. 
and 755 mm. The 
levels in B and D are 
made the same (by 
raising one bottle) — 
to put the gas under 
atmospheric pres¬ 
sure. The water in 
D is measured; its 
volume is the same 
as the volume of the 
oxygen, say 1.75 li¬ 
ters. AF is weighed; 
its loss is the weight of the oxygen, say 2.322 gm. (c) The observed 
volume is reduced to standard conditions; thus: 1.75 X X yM = 



Fig. 


157. — Apparatus for finding the weight of 
a liter of oxygen. 


1.625. (d) The weight of one hter of oxygen at 0° C. and 760 mm. is 

found by dividing the weight of the oxygen by its corrected volume; 
thus, 2.322 H- 1.625 = 1.429. 

We shall see later (§ 348, last paragraph) that allowance for the pres¬ 
sure of water vapor must be made in the numerical value of the final 
pressure when a gas is collected over water. 


EXERCISES 

1. State Charles’s law. Illustrate it. 

2. State Boyle’s law. Illustrate it. 

3. Give examples from everyday life of {a) expansion and of (6) con¬ 
traction of gases caused by change of temperature. 

4. Apply Exercise 3 to change of pressure. 

6. (a) Change these centigrade readings to absolute: 100, 0, 15, 
— 15, 250, 273, — 273. (6) Change these absolute readings to centi¬ 

grade: 273, 0, 200, 100, 473, 180, 373. 


PROBLEMS 

1. Reduce the following to the volume occupied at 760 mm.: (a) 20 
cc. at 745 mm.; (6) 45 cc. at 765 mm.; (c) 450 cc. at 755 mm.; (d)1.51. 
at 763 mm. 

2. Reduce the following to the volume occupied at 0° C.: (a) 170 cc. 
at 80° C.; (6) 450 cc. at 15° C.; (c) 70.6 cc. at 17° C.; (d) 49 cc. at 
19° C. 




















LAWS AND THEORIES 


271 


3. Reduce the following to the volume at standard conditions: 
(a) 250 cc. at 780 mm. and 20° C.; (6) 140 cc. at 745 mm. and 21° C. 

4. Reduce the following to the volume at standard conditions: 
(a) 247 cc. at 720 mm. and 14° C.; (6) 1000 cc. at 750 mm. and 18° C. 


SUGGESTIONS FOR LABORATORY WORK 
(References are to Newell's Laboratory Exercises in Chemistry) 
Exercise 41 — Weight of 22.4 Liters of Oxygen. 


TOPIC VII: LAWS AND THEORIES — KINETIC- 
MOLECULAR THEORY 

346. Chemical laws and theories. — Many related facts 
are often summarized in a concise statement called a law. 
The explanation of a group of related facts is called a theory. 

347. The kinetic-molecular theory of the structure of 
gases. — Gases behave alike in many ways. (1) They 
expand indefinitely and fill the whole containing vessel. 
This tendency is shown by the pressure exerted on the inside 
walls of the vessel or by the rapid escape through a hole, e.g., 
a puncture in an automobile tire. (2) They diffuse readily, 
^.c., they mix quickly and uniformly, and likewise pass 
through porous substances. Thus, ammonia is quickly 
detected in a room when a bottle of this substance is opened. 
(3) They are highly compressible, e.g., a large volume of 
air can be pumped into an automobile tire; recall also 
that a large volume of oxygen, hydrogen, and carbon dioxide 
can be forced into small metal cylinders. (4) Gases change 
uniformly in volume with changes in temperature and in 
pressmre. 

The uniform behavior of gases led to a theory of their struc¬ 
ture called the kinetic-molecular theory. The essentials of 
this theory are as follows: — 

(1) A gas is made up of a large number of exceedingly 
minute particles called molecules. The molecules of a 
given gas are alike. 

(2) The molecules of a gas are relatively far apart, i.e., the 
distances between molecules are large compared with the 


272 


A BRIEF COURSE IN CHEMISTRY 


molecules. They are exceedingly small particles scattered 
throughout a very large space. 

(3) The molecules of a gas are moving rapidly in all direc¬ 
tions, striking each other and the walls of the containing 
vessel. Being elastic, they rebound and continue to move 
ceaselessly without loss of velocity. These blows, or impacts, 
produce pressure, and the series of rapid blows produces the 
effect of uniform and constant pressure. 

(4) The movements of the molecules of a gas increase in 
velocity as temperature rises, and decrease as it falls. A 
definite change in temperature changes the average velocity 
of the molecules, thereby producing a definite change in 
volume (if the pressure is kept constant). 

(5) The molecules have little or no tendency to cohere, 
^.6., they lead an independent existence, being far apart 
relatively, except at great concentration. 

Let us interpret the behavior of gases by the kinetic- 
molecular theory. 

Consider the law of Boyle. When a gas is compressed, 
the molecules are pressed more closely together. And when 
a gas is compressed to half its volume, the pressure produced 
by the blows of the moving molecules is doubled because the 
number of blows per second against the walls is doubled. 

Consider the law of Charles. The molecules move more 
rapidly, and consequently require more space when the tem¬ 
perature is raised; and the change in the average velocity 
of the molecules is proportional to the change in temperature, 
^.e., the expansion (or contraction) is uniform with the tem¬ 
perature. 

Again, the rapid movement of gases called diffusion is ade¬ 
quately interpreted by the constant motion of the molecules. 

Finally liquefaction, which occurs when a gas is reduced 
to a low temperature and subjected to great pressure (§ 168), 
is explained as follows: the lowering of temperature slows 
up the velocity of the molecules and the increasing of pressure 
forces them nearer together, so that their fundamental tend¬ 
ency to cohere slightly becomes greater and manifests itself 
in the formation of larger particles (than molecules) which 
eventually become a mass of liquid. 


VAPOR PRESSURE 


273 


TOPIC VIII: VAPOR PRESSURE — VOLUMETRIC COM¬ 
POSITION OF WATER — GAY-LUSSAC’S LAW OF 
GAS VOLUMES — HYDROGEN PEROXIDE 

348. Vapor pressure. — Evaporation produces pressure, 
which is called the vapor pressure of water, or merely vapor 
pressure. 

The fact that evaporation produces pressure can be shown 
by a simple experiment. 

The dry bottle (Fig. 158, left) is fitted with an open U-shaped tube 
partly filled with a colored liquid {e.g., dilute ink) to serve as an indicator 
of pressure. When 
the stopper is re¬ 
moved, a httle water 
poured into the bot¬ 
tle, and the stopper 
quickly replaced, the 
water begins to 
evaporate. As the 
water evaporates, 
the colored liquid 
shows an increase in 
pressure (Fig. 158, 
right). 

The amount of — Experiment to illustrate vapor pressure, 

pressure exerted In the dry bottle (on the left) the colored liquid in 
by water vapor bent tube shows no pressure. Whereas, after 

1 . 1 water has been added, the colored liquid shows 

aepenas solely on pressure in the bottle (on the right), 
the temperature 

— not on the quantity of water. This is readily seen by 
comparing the heights of the mercury in the fine-bore tubes 
shown in Fig. 159. 

Each tube was first filled with mercury and inverted in the dish of 
mercury. In each tube the mercury sank to the same point (760 mm.). 
In tube A there is no water vapor in the space above the mercury, and 
the height of the mercury column is 760 mm. A drop of water was 
forced up into the tubes B and C by means of a medicine dropper. In 
B the space above the mercury is filled with water vapor kept at 20° C.; 
the vapor exerts a pressure and forces the mercury down to nearly 742 
mm. That is, water vapor at 20° C. exerts a pressure equal to 18 mm. of 



















274 


A BRIEF COURSE IN CHEMISTRY 


mercury. Similarly, in C the space is filled with water vapor kept at 
50° C., and the mercury is forced down to 668 mm., the water vapor at 
this temperature exerting a pressure of 92 mm. 


Vapor pressure and boiling point are related. The vapor 
escaping from water boiling in an open vessel overcomes the 
pressure of the atmosphere upon the surface of the water. 
Since the normal pressure of the atmosphere is 760 mm., the 
normal boiling point, so to speak, is 100° C. The boiling 
point becomes lower as the pressure is de¬ 
creased and higher as the pressure is in¬ 
creased. Warm water will boil under the 
receiver of an air pump and on the top of a 
high mountain, because under these condi¬ 
tions the pressure is less than normal. In a 
pressure cooker the temperature of the water 
is above 100° C. On the other hand in a 
vacuum vessel, such as is used to evaporate 
sugar solutions, the boiling point is often as 
low as 70° C. 

The pressure exerted by water vapor has 
a definite value for each temperature. These 
values can be found in the Table of Vapor 
Pressure given in the Appendix. 

A practical application of vapor’ pressure is often 
made in computing the weight of a gas measured 
over water. Thus, in finding the weight of a liter of 
oxygen (§ 346) the gas is allowed to stand over the 
water long enough to become saturated with water 
vapor, that is, the bottle finally contains a mixture 
of oxygen and the maximum amount of water vapor 
at the given temperature. In such a mixture, each gas (oxygen and 
water vapor) shares the total atmospheric pressure. Hence the actual 
pressure exerted by the oxygen is found by subtracting the pressure of 
the water vapor (a in the table of vapor pressure) from the total pres¬ 
sure (read on the barometer). The corrected pressure is used in the 
formula for reducing the volume of a gas to its volume at 0° C. and 
760 mm. (§ 346, next to last paragraph). Thus, if the temperature is 
20° C., the vapor pressure as found in the table is 17.4. The formula 
then becomes — 


B 


Fig. 159. — Ex¬ 
periment to 
show that the 
amount of 
vapor pressure 
depends on the 
temperature. 


^ 7,x 273 ^(755 -iM = , 

293 760 


760 














VAPOR PRESSURE 


275 


349. Efflorescence and vapor pressure. — The water vapor 
escaping, however slowly, by the efflorescence of a hydrate 
(crumbling of a crystallized salt due to loss of water of 
hydration, § 86), exerts a slight vapor pressure. If this vapor 
pressure is greater than the pressure of the water vapor in 
the atmosphere, the substance will lose water of hydration. 
That is, the substance will effloresce. Some hydrates do not 


effloresce at all, or only slightly, because their vapor pressure 
is less than the average vapor pressure of the air. Crystal¬ 
lized barium chloride (BaCb . 2 H 2 O) and 
gypsum (selenite — CaS 04.2 H 2 O) belong to 
the latter class. 

350. Deliquescence and vapor pressure. — 

Water vapor from the air condenses on the 
surface of soluble solids and produces a very 
concentrated solution. Such a solution has a 
much lower vapor pressure than the average 
pressure of the water vapor in the air. The 
solution, therefore, continues to take up water 
until its vapor pressure equals the pressure of 
the water vapor in the air. 

351. Volumetric composition of water. — 

The volumetric composition of water is de¬ 
termined by measuring the volumes of hydrogen 
and oxygen that combine to form water. 



Fig. 160.-Ap¬ 
paratus for 
determining 
the volumet- 
ric com¬ 
position of 
water. 


A sketch of one form of apparatus is shown in Fig. 160. 

The essential part is the eudiometer A. It is a gradu¬ 
ated glass tube closed at the upper end. Near this end 
two platinum wires are fused into the glass. The outer ends are looped 
and the inner ends are near together so that an electric spark will leap 
across the gap and produce enough heat to cause the two gases to unite. 

The eudiometer is filled with mercury and inverted in the jar of the 
same liquid. Hydrogen is introduced until the eudiometer is about one 
fourth full. The mercury levels inside and outside are made the same, 
and the volume of hydrogen is read accurately. The temperature and 
pressure are also read. Approximately an equal volume of oxygen is 
then introduced, the levels are adjusted, and the total volume read accu¬ 
rately. Each volume is corrected for temperature and pressure (§ 344), 
The difference between the two corrected volumes is the volume of oxy¬ 
gen. An excess of oxygen is needed to lessen the violence of the explo¬ 
sion ; this excess takes no part in the chemical change. 










276 


A BRIEF COURSE IN CHEMISTRY 




The combination of the two gases is caused by connecting the looped 
ends of the platinum wires with an induction coil and battery, and pass¬ 
ing a spark across the gap. A slight explosion indicates combination. 
The mercury, after the shock from the explosion, rises and nearly fills the 
eudiometer. The volume of the water pro¬ 
duced is too minute to measure in this appa¬ 
ratus. After the mercury and residual gas 
(oxygen) are cool, the levels are adjusted, 
and the volume of gas is read (as well as the 
temperature and pressure). The corrected 
volume of this gas is subtracted from the 
corrected volume of oxygen, thus giving the 
actual volume of oxygen that combined with 
all the hydrogen to form water. 

An example will make this experiment 
clear. Suppose the corrected volumes 
were: — 

Volume of hydrogen added . 22.3 cc. 

Volume of hydrogen and oxy¬ 
gen added.41.5 cc. 

Volume of oxygen added . . 19.2 cc. 

Volume of oxygen left . . . 8.0 cc. 



This means that 19.2 — 8, or 11.2, cc. of 
oxygen were actually used. In other words 
the two gases combined in the ratio of 22.3 
to 11.2, or very nearly 2 volumes of hydro¬ 
gen to 1 volume of oxygen. 

362. Gay-Lussac^s law of gas vol¬ 
umes. — A modification of the experi¬ 
ment described in § 361 shows that 2 
volumes of water vapor are formed 
when 2 volumes of hydrogen and 1 
volume of oxygen unite. 


In the apparatus (Fig. 161) the eudiometer 
is surrounded by a large tube through which 
steam is passed, thereby preventing the con¬ 
densation of the water vapor. If 2 volumes 
of hydrogen and 1 volume of oxygen are 
exploded, 2 volumes of water vapor are 
formed — provided all the gases are measured at the same pressure and 
the same temperatures (about 100° C). This result means that the 
volumes of the three gases — hydrogen, oxygen, and steam — involved 
in this reaction are expressed by 2, 1, and 2, i.e., by small whole 
numbers. 


Fig. 161.—Apparatus for 
showing that 2 volumes 
of water vapor are 
formed by exploding 2 
volumes of hydrogen 
and 1 volume of oxygen. 





















GAY-LUSSAC’S LAW 


277 



When other chemical changes involving gases are studied, 
simple numerical relations are also found. This general 
fact, which was discov¬ 
ered by the French chem¬ 
ist Gay-Lussac (Fig. 162), 
in 1808, may be stated as 
a law, thus: — 

In a chemical change the 
volumes of the gases can he 
expressed hy small whole 
numbers. 


This law will be used 
in a later topic (§ 363). 


EXERCISES 


1. Describe the experiment 
which shows (a) evaporation 
produces vapor pressure and 
(6) the amount of vapor pres¬ 
sure depends on temperature. 

2. What is the vapor pres¬ 
sure of water at 20° C. ? At 
100° C.? 


Fig. 162. — The French chemist Gay- 
Lussac (1787-1850), who first studied 
the simple relation of gas volumes in 
chemical changes. 


3. Interpret {a) efflores¬ 
cence and (5) deliquescence by the principle of vapor pressure. 

4. State the volumetric composition of water. How is it found? 
6. State and illustrate Gay-Lussac’s law. 


PROBLEMS 

1. What would be the volume of the dry gas at 0° C. and 760 mm. ? 
(a) 80 cc. at 750 mm. and 17° C.; (b) 100 cc. at 765 mm. and 19.5° C. 

2. A mixture of 500 cc. of oxygen and 1250 cc. of hydrogen — both 
at the normal temperature and pressure — is exploded. What volume, 
if any, of gas will remain? How much? Which gas? 

353. Hydrogen peroxide. — This is a compound of hy¬ 
drogen and oxygen, which decomposes readily into oxygen 
and water. The commercial solution (3 per cent) is used 
to bleach hair, etc., and to sterilize wounds. 





278 


A BRIEF COURSE IN CHEMISTRY 


TOPIC IX: LAWS OF THE CONSERVATION OF MATTER 

AND DEFINITE PROPORTIONS — ATOMIC THEORY — 
ATOMS AND MOLECULES — ATOMIC WEIGHTS 

354. Law of the conservation of matter. — The terms 
law and theory are defined in § 346. 

In a chemical change the total weight of the matter involved 

is not altered. That is, the 
sum of the weights of the origi¬ 
nal substances equals the sum 
of the weights of the final 
substances. This general 
fact about chemical change is 
summed up by the law of the 
conservation of matter, thus:— 

No weight is lost or gained 
in a chemical change. 

355. Law of constant com¬ 
position. — In all chemical 
compounds the different con¬ 
stituents are present in a 
definite and constant propor¬ 
tion by weight. This general 
fact, stated in the form of a 
law, becomes the law of con¬ 
stant composition or the law of definite proportions, thus: — 

A given chemical compound always contains the same ele¬ 
ments in the same proportion by weight, 

356. The atomic theory. — The theory that interprets 
the facts summarized in the two fundamental laws just stated 
is called the atomic theory. It was proposed about 1803 
by the English chemist Dalton (Fig. 163). The essential 
points in Dalton’s atomic theory are: — 

(1) Elements are made up of very small particles called 
atoms. 

(2) Atoms of the same element are alike and have an un¬ 
varying weight, called the atomic weight. 



Fig, 163. — The English chemist 
Dalton (1766-1844), who proposed 
the atomic theory. 





CONSERVATION OF MATTER 279 

(3) Atoms of different elements differ from one another 
in weight. 

(4) In a chemical change undivided atoms unite, are 
separated, or exchange places in the ratio of their particular 
weights. 

The atomic theory means in a few words that matter is 
composed of atoms, which remain undivided in chemical 
changes. 

With the help of the atomic theory, many facts about 
elements, compounds, and chemical change can be made 
much clearer (§§ 358, 369). 

367. Atoms and molecules. — An atom is the smallest, 
indivisible particle of an element. It is the smallest part of 
an element which participates in a chemical change. 

A molecule is a chemical combination of two or more 
atoms. If the atoms are alike, the molecule is a molecule of 
an element, e.g., a molecule of oxygen consists of two atoms 
of the element oxygen (O 2 ). If the atoms are not alike, the 
molecule is a molecule of a compound, e.g., a molecule of 
water consists of two atoms of hydrogen and one of oxygen 
(H 2 O). A molecule is the smallest particle of a compound. 
Thus, a molecule of water is the smallest particle of the com¬ 
pound water, because if we try to produce a smaller particle, 
we obtain atoms of the elements hydrogen and oxygen. 

According to present views an atom is essentially electri¬ 
cal. It consists of a nucleus surrounded by electrons 
moving in orbits. The nucleus is positively charged. 
Each electron is negatively charged. The positive charge 
on the nucleus is electrically balanced by the total negative 
charges on the surrounding electrons. Hence an atom as 
a whole is electrically neutral. A molecule is a group of 
atoms held together electrically. (Refer to §§ 358, 424,529.) 

368. Interpretation of chemical change by the atomic 
theory. — (1) A sample of the element copper consists of 
many millions of atoms of copper. A sample of oxygen like¬ 
wise consists of atoms of oxygen. When the chemical 
change occurs between copper and oxygen, atoms of copper 
combine with atoms of oxygen and form molecules of the 
compound copper oxide. And this combining of atoms into 


280 


A BRIEF COURSE IN CHEMISTRY 


molecules continues until all the atoms of copper or all the 
atoms of oxygen (or under certain conditions all the atoms of 
both elements) are used up. 

(2) Similarly, the decomposition of mercuric oxide is the 
separation of numberless molecules of the compound mer¬ 
curic oxide into atoms of the elements mercury and oxygen. 

(3) So also, the liberation of hydrogen from sulfuric acid 
is the exchange of atoms of zinc for the atoms of hydrogen in 
the compound sulfuric acid. 

Chemical change can also be interpreted by the electron 
theory. Atoms lose or gain electrons more or less readily. 
Thus, copper atoms tend to lose electrons and oxygen atoms 
to gain them. When atoms of copper and of oxygen form 
molecules of copper oxide, each copper atom loses two 
electrons which are gained by an oxygen atom. The copper 
atoms thereby become positive and the oxygen atoms nega¬ 
tive. And each pair of oppositely charged atoms is held 
together electrically as a molecule of copper oxide. 

359. How the atomic theory helps us explain the two fun¬ 
damental laws of chemical change. — (1) The law of the 
conservation of matter states that there is no loss or gain of 
weight in a chemical change. According to the atomic the¬ 
ory the weight of an atom is never changed. This means 
that in the case of copper and oxygen, no atoms are created 
or destroyed. Hence the weight of the copper oxide formed 
must equal the sum of the weights of the copper and oxygen 
used up chemically. Inasmuch as in all chemical changes 
there is no loss or gain in weight, it is obvious that the atomic 
theory, which assumes unvarying weights of atoms, is in 
accord with the law of the conservation of matter. 

(2) The law of constant composition states that a given 
compound always consists of the same elements in a fixed 
proportion. According to the atomic theory, molecules 
are formed by the union of a definite number of whole atoms 
of each element. Each molecule of the compound copper 
oxide, for example, would consist of one or more atoms of 
copper united with one or more atoms of oxygen. Therefore 
the composition of each molecule of copper oxide would be 
constant, ^.e., each molecule would consist of the same ele- 


CONSERVATION OF MATTER 


281 


merits united in a constant proportion by weight. This 
means that the composition of copper oxide would always be 
a certain per cent of copper and a certain per cent of oxygen. 
Since all chemical compounds have been found to have a 
constant composition, the atomic theory is in harmony with 
the law of constant composition. 

360. Atomic weights. — According to the atomic theory 
atoms of the same element always have the same weight but 
atoms of different elements have different weights. This 
means, for example, (1) that an atom of oxygen throughout 
all its varied changes retains its weight, and (2) that this 
weight differs from the weight of other kinds of atoms. 
The weights of different kinds of atoms are called the atomic 
weights of the elements, or briefly atomic weights. These 
weights have been determined by very accurate experiments 
(§ 377). A table giving the exact, as well as the approxi¬ 
mate, values can be found on the inside of the back cover of 
this book. The standard atomic weight is oxygen = 16. 

Atomic weights are relative weights. That is, the atomic 
weight of copper is 63.57, not 63.57 gm. or any other actual 
weight, but 63.57 as long as 16 is accepted as the standard 
atomic weight of oxygen. 

The exact determination of atomic weights is a difficult 
task. Until this subject is discussed (§§ 376-378), it will 
be well enough to use the approximate weights from the 
table as needed. 


EXERCISES 

1. Define law and theory as used in science. 

2. State the law of conservation of matter. Illustrate it by the 
decomposition of mercuric oxide. 

3. State the law of constant composition. Illustrate it by the 
gravimetric composition of water. 

4. State the four points in the atomic theory. 

6. (a) What is an atom ? A molecule ? (6) Discuss the relation 

of atoms to molecules. 

6 . Describe a chemical change by the (a) atomic and (6) electron 
theory. 

7. Interpret by the atomic theory the two laws: (a) conservation 
of matter, and (6) constant composition. 

8 . What are atomic weights? 


282 


A BRIEF COURSE IN CHEMISTRY 


9. Learn the approximate atomic weight of these elements: Hydro¬ 
gen, oxygen, nitrogen, carbon, copper, iron, sulfur, chlorine. 

10. As in Exercise 9: Sodium calcium, lead, magnesium, mercury, 
silver, zinc. 

SUGGESTIONS FOR LABORATORY WORK 
(References are to Newell’s Laboratory Exercises in Chemistry) 
Exercise 22 — Formation of the Compound Copper Sulfide. 


TOPIC X: EQUIVALENT WEIGHTS 

361. Equivalent weights. — The equivalent weight of an 

element is the weight that combines with or replaces 8 grams 
of oxygen. For example, 12 grams of magnesium combines 
with 8 grams of oxygen. Therefore 12 is the equivalent 
weight of magnesium. Similarly, hydrogen and oxygen 
combine in the ratio of 1 to 8; therefore 1 is the equivalent 
weight of hydrogen. 

Equivalent weights are sometimes called reacting weights, 
or simply equivalents. The term equivalent weights is pref¬ 
erable, because they actually are the weights chemically 
equivalent to each other. Thus, 1 gm. of hydrogen com¬ 
bines with 35.5 gm. of chlorine, and this 1 gm. of hydrogen 
can be replaced chemically by 32.5 gm. of zinc, 12 gm. of 
magnesium, 9 gm. of aluminum, or 20 gm. of calcium, and 
so on. These elements are chemically equivalent in the 
ratio of these weights. 

In the laboratory the equivalent weights of certain ele¬ 
ments can not be determined directly in terms of oxygen. 
They are determined by finding the weight that combines 
with or replaces the equivalent weight of some other element, 
usually hydrogen. Thus, 32.5 grams of zinc replace 1 gram 
of hydrogen from hydrochloric acid; therefore 32.5 is the 
equivalent weight of zinc. 

Comparison of the equivalent weight with the (approxi¬ 
mate) atomic weight of the same element reveals an impor¬ 
tant relation, as may be seen by Table XI. An examination 
of this comparative table shows an integral relation between 
equivalent weights and atomic weights. In other words. 


EQUIVALENT WEIGHTS 


283 


TABLE XI. — Comparison of Equivalent Weights 
AND Atomic Weights 


Element 

Equivalent 

Weight 

Atomic Weight 

Multiple 

Oxygen. 

8 

16 

2 

Aluminum .... 

9 

27 

3 

Bromine .... 

80 

80 

1 

Calcium .... 

20 

40 

2 

Carbon . 

3 

12 

4 

Chlorine .... 

35.5 

35.5 

1 

Hydrogen . . . . 

1 

1 

1 

Magnesium . . . 

12 

24 

2 

Silver. 

108 

108 

1 

Sodium. 

23 

23 

1 

Sulfur. 

16 

32 

2 

Zinc. 

32.5 

65 

2 


the atomic weight of an element is identical with its equiva¬ 
lent weight or is a simple integral multiple of it. These 
integral multiples are the same as the numerical valence 
of the element (§ 117). 

362. The difference between equivalent and atomic 
weights. — The equivalent weight of an element, as pre¬ 
viously stated, is the weight that combines with or displaces 
8 gm. of oxygen. That is, equivalent weights are merely 
weights in a scheme which is based on O = 8. But 8 is not 
the unit used in chemistry to express the composition of 
compounds or the quantitative relations of the elements. 
The quantitative unit is 16. Equivalent weights are em¬ 
pirical. Whereas atomic weights are part of the whole 
structure of theoretical chemistry which is erected on all 
the facts, laws, and theories of chemistry. 

EXERCISES 

1 . Define and illustrate equivalent weight. 

2 . What is the equivalent weight of oxygen, hydrogen, zinc, mag¬ 
nesium, aluminum, sodium, calcium, chlorine, silver? 

3. What is the difference between an equivalent weight and an atomic 
weight ? 














284 


A BRIEF COURSE IN CHEMISTRY 


PROBLEMS 

1. Calculate the equivalent weights of the respective metals from 
the following data: (a) 0.5 gm. of calcium unites with 0.2 gm. of oxygen 
to form calcium oxide (CaO). (b) 15 gm. of mercury unite with 1.2 gm. 
of oxygen to form mercuric oxide (HgO). 

2. Calculate the equivalent weight of sodium from the following: 
(a) 2.3 gm. of sodium liberate 0.1 gm. of hydrogen from water, (b) 1.15 
gm. of sodium hberate 555.5 cc. of hydrogen (at standard conditions). 

SUGGESTIONS FOR LABORATORY WORK 
(References are to Newell’s Laboratory Exercises in Chemistry) 

Exercise *21 — Equivalent Weight of Zinc. 

Exercise 20 — Equivalent Weight of Magnesium. 

Exercise SI8 — Equivalent Weight of Aluminum. 


TOPIC XI: GAY-LUSSAC’S LAW — AVOGADRO’S LAW —MO¬ 
LECULAR WEIGHTS — MOLECULES — FORMULAS — 
MOLECULAR EQUATIONS 

363. Gay-Lussac’s law of gas volumes. — This law was 
discussed in § 362. 

The volumes of gases involved in several chemical changes 
are shown in Table XII. 

I 

TABLE XII. — Combination of Gases by Volume 


Volumes of Combining Gases 

Volumes of Gaseous Product 

2 volumes of hydrogen 

1 volume of oxygen 

2 volumes of water vapor 

1 volume of hydrogen 

1 volume of chlorine 

2 volumes of hydrogen chloride 

1 volume of nitrogen 

3 volumes of hydrogen 

2 volumes of ammonia 

1 volume of nitrogen 

1 volume of oxygen 

2 volumes of nitric oxide 








GAY-LUSSAC’S LAW 


285 


It is clear from the above table that the volumes of the 
gases can be expressed by small whole numbers. This 
simple relation is true of all gas reactions, and may be stated 
as Gay-Lussac’s law, thus: — 

In a chemical change the volumes of the gases can be ex¬ 
pressed by small whole numbers. 

This law applies only to gases. For example, carbon is not 
a gas, though it is frequently involved in reactions with 
gases, e.g.: — 

Carbon + Oxygen = Carbon Dioxide 

1 vol. 1 vol. 

So we omit carbon as far as volume is concerned, and say 
carbon unites with 1 volume of oxygen (O 2 ) to form 1 volume 
of carbon dioxide (CO 2 ). 

364. Avogadro’s law. — In 1811 the Italian physicist 
Avogadro proposed an explanation of the simple numerical 
relation of gas volumes in a chemical change. It is usually 
called Avogadro’s law and may be stated thus: — 

Equal volumes of gases under like conditions of temperature 
and pressure contain the same number of molecules. 

This law means, for example, that a liter of oxygen con¬ 
tains just as many molecules as a liter of hydrogen, nitric 
oxide, or any other gas, if the temperature and pressure are 
the same. Hence, if equal volumes of gases are weighed, the 
weights are in the same ratio as the weights of single mole¬ 
cules of the gases. 

365. How Avogadro’s law is used to find molecular weights. 

— Let us consider carbon dioxide and oxygen. A liter of 
carbon dioxide weighs 1.98 gm. and a liter of oxygen 1.43 gm. 
at 0° C. and 760 mm. Therefore the weight of a liter of 
carbon dioxide is 1.38 times that of a liter of oxygen {i.e., 
1.98 -i- 1.43 = 1.38). Since a liter of each gas contains 
the same number of molecules, the weight of the carbon 
dioxide molecules is 1.38 times the weight of the oxygen 
molecules. This means that the molecular weight of carbon 
dioxide is 1.38 times the molecular weight of oxygen. 

The molecular weight of oxygen is 32 (see next paragraph). 
Hence the molecular weight of carbon dioxide is 32 times 


286 


A BRIEF COURSE IN CHEMISTRY 


1.38 = 44. Therefore 44 is the molecular weight of carbon 
dioxide. 

366. Why the molecular weight of oxygen is 32. — There 

are two reasons. (1) An atom of oxygen weighs 16, because 
this number has been adopted by chemists as the standard 
atomic weight (§ 378). (2) A molecule of oxygen contains 

2 atoms, as will be shown in § 368. Hence the molecular 
weight of oxygen is 32 (i.e., 2 times 16). 

367. Finding molecular weights by the vapor density 
method. — The method of finding molecular weights de¬ 
scribed in § 365 is called the vapor density method. The 
steps are (1) find the. vapor density (i.6., relative weight) 
referred to oxygen, and (2) multiply this value by 32. 

The expression ‘‘ vapor density referred to oxygen ” means 
the number found by dividing the weight of a given volume 
of a gas or vapor by the weight of an equal volume of oxygen 
(measured at the same temperature and pressure). Thus, 
in the example given above the number 1.38 is the vapor 
density of carbon dioxide (i.e., 1.98 -i- 1.43). 

Therefore, in brief: — 

Molecular Weight = Vapor Density referred to Oxygen 
times 32. 

Some substances can not be vaporized without decomposition. The 
molecular weights of such substances can not, of course, be found by 
the vapor density method. 

No experimental method has been devised for determining the molec¬ 
ular weight of a substance in the solid state {i.e., not dissolved or vapor¬ 
ized) ; it is customary to assume that the molecular weight of such 
substances is the sum of the atomic weights in the simplest formula 

(§§ 104, 106). 

368. A molecule of oxygen contains two atoms. — We 

use Gay-Lussac’s and Avogadro’s laws to prove that a mole¬ 
cule of oxygen contains two atoms. When oxygen and 
nitrogen combine to form nitric oxide, the volumes used 
and produced can be expressed (according to Gay-Lussac’s 
law) thus: — 

Oxygen + Nitrogen = Nitric Oxide 

1 vol. 1 vol. 2 vols. 

Now according to Avogadro’s law, equal volumes of gases 


GAY-LUSSAC’S LAW 


287 


contain the same number of molecules. Suppose there are 
1000 molecules of oxygen in 1 volume. Then by Avogadro’s 
law there are 2000 molecules of nitric oxide in 2 volumes 
of this gas. Now every molecule of nitric oxide must con¬ 
tain at least one atom of oxygen; and the 2000 molecules 
must contain at least 2000 atoms of oxygen. But these 
2000 atoms of oxygen were provided by the 1000 molecules 
of oxygen. Therefore, each molecule of oxygen must con¬ 
tain at least 2 atoms of oxygen. 

369. Molecules of other elementary gases contain two 
atoms. — There are good reasons for dropping the “ at least ” 
and saying a molecule of oxygen contains 2 atoms. Thus, 
there is no reaction in which a given volume of oxygen and 
other elementary gases, e.g., nitrogen, hydrogen, and chlorine, 
provides material for more than two volumes of the gaseous 
product. This means there is no reaction in which a mole¬ 
cule of these gases is divided into more than two parts. And 
so we conclude that the molecule of these gases contains 
only two atoms. Hence we write their formulas O 2 , N 2 , H 2 , 
and CI 2 . It is clear now why we used these formulas in 
preceding sections for molecules of the gases. 

370. A mole. — A mole of a substance is the number of 
grams numerically equal to its molecular weight. It is 
sometimes called a gram-molecular weight. Thus, a mole 
of carbon dioxide is 44 grams because the molecular weight 
is the number 44. Similarly, a’ mole of oxygen is 32 grams, 
of carbon monoxide is 28 grams, and of nitric oxide is 30 
grams. 

The volume occupied by one mole of oxygen (at 0° C. and 
760 mm.) is 22.4 liters. Suppose we construct a box holding 
22.4 liters (Fig. 164), and fill it with oxygen (at 0° C. and 
760 mm.), the gas will weigh 32 grams. This must be so, 
because one liter of oxygen weighs 1.43 grams, and the volume 
occupied by 32 grams of oxygen will be 32 1.43, or 22.4 

liters (in round numbers). Similarly, the volume occupied 
by one mole of carbon monoxide is 22.4 1., i.e., 28 1.25 

= 22.4. 

This volume (22.4 liters) is sometimes called the gram- 
molecular volume, since it is the volume of the gram-molec- 


288 A BRIEF COURSE IN CHEMISTRY 

ular weight. It is also called a uni-molar volume, because 
it is the volume of one mole. So it is clear that 1 mole of a 
gas and 22.4 liters of the same gas are equal — one being 

the weight and the other the 
volume of the same mass of 
gas. 

371. How molecular 
weight is calculated by the 
molar method. — Suppose 
we construct several cubical 
boxes each holding 22.4 liters 
(Fig. 165), weigh each box, 
fill each with a gas (at 0° C. 
and 760 mm.), and weigh 
again. In each case the 
weight of the several gases 
is numerically equal to the molecular weight. If the gases 
are nitric oxide, carbon monoxide, hydrogen chloride, and 
ammonia, the numbers obtained are 30, 28, 36.5, and 17. 
And these numbers are the molecular weights respectively 
of these gases. 

Hence to find the molecular weight of a gas by the molar 
method, calculate the weight of 22.4 liters of the gas. . It 



Fig. 164.—A box holding 1 mole or 
22.4 liters of a gas (actual length of 
one edge is 28.2 cm.). 



Fig. 165. — A mole of different gases occupies 22.4 liters. A = 30 gm. NO, 
R = 28 gm. CO, C = 36.5 gm. HCl, D = 17 gm. NH3. 

is not necessary to use such a large volume in the actual 
experiment. We simply find the weight of any convenient 
volume (from a table or by experiment) and then calculate 
the weight of 22.4 liters. Thus, if 1.5 liters of carbon mon¬ 
oxide (at 0° C. and 760 mm.) weigh 1.88 grams, the weight 
of 22.4 liters is found thus: 1.88 ^ 1.5 = 1.25; 1.25 X 22.4 
= 28. Therefore 28 is the molecular weight of carbon 
monoxide. 

















GAY-LUSSAC’S LAW 


289 


372. Molecular formulas of compounds. — A formula 
of a compound represents by symbols the kind and number 
of atoms in a molecule. 

In § 105 it was shown that the simplest formula of a com¬ 
pound can be calculated from the percentage composition 
by dividing the per cent of each element in the compound 
by the respective atomic weight, and then, if necessary, 
reducing the quotients to the smallest whole numbers. A 
formula thus calculated is called the simplest formula be¬ 
cause it expresses in the simplest chemical way the propor¬ 
tions of the different elements in a compound. But it may 
not be its correct formula. 

The correct formula of a compound must represent its 
molecular weight. That is, the sum of the weights repre¬ 
sented by the kind and number of atoms in the correct 
formula must be equal (or very nearly equal) to the molec¬ 
ular weight found by experiment. Let us take an example. 
A compound was found by analysis to contain 92.3 per cent 
of carbon and 7.7 of hydrogen, and to have a vapor density 
of 2.4375. Dividing the percentages by the atomic weights, 
we have: 92.3 12 = 7.7, and 7.7 1 = 7.7. Since 

7.7 :7.7 as 1 :1, the compound contains at least one atom 
each of carbon and hydrogen, and its simplest formula is 
CH. This formula corresponds to the molecular weight 13. 
But the vapor density 2.4375 requires the molecular weight 
78 (Le., 2.4375 times 32 = 78), which is six times the weight 
(13) corresponding to the formula CH. Hence the molecular 
formula of this compound is not CH, but CeHe. 

If the molecular weight of a compound can not be found 
by experiment, then the simplest formula is accepted as 
the molecular formula. For example, a certain compound 
contains 40 per cent of calcium, 12 of carbon, and 48 of 
oxygen. Dividing each per cent by the proper atomic weight, 
we have: 40 40 = 1, 12 12 = 1, 48 16 = 3. That 

is, one molecule of this compound contains (at least) 1 atom 
each of calcium and carbon, and 3 of oxygen; therefore the 
simplest formula is CaCOs. This is also accepted as its 
molecular formula, because the molecular weight can not be 
found by any known method. 


290 


A BRIEF COURSE IN CHEMISTRY 


373. Molecular formulas of elements. — Several gaseous 
elements have molecular weights which are twice the atomic 
weight. This means that the molecule consists of two atoms, 
and their molecular formulas are, for example, O2, H2, CI2, N2 
(also Br2). (Compare § 369.) 

374. Molecular equations. — Reactions involving the 
common elementary gases should be expressed by molecular 
equations. Thus, the molecular equation for the formation 
of water vapor from hydrogen and oxygen is: — 

2 H2 + O2 = 2 H2O 

This equation is read: Two molecules of hydrogen unite 
with one molecule of oxygen to form two molecules of water 
vapor. Since this equation correctly represents the interact¬ 
ing substances as molecules, the equation is correctly called 
a molecular equation. A molecular equation is sometimes 
called a volumetric equation or a gas equation, because it 
shows the volumes of gases involved in the reaction. Thus, 
the equation 

H2 + CI2 = 2 HC 1 

1 molecule 1 molecule 2 molecules 

may be written : — 

H2 + CI2 = 2 HC 1 

1 volume 1 volume 2 volumes 

because equal numbers of molecules represent equal volumes. 
The second equation is read : One volume of hydrogen and one 
volume of chlorine form two volumes of hydrogen chloride. 
It should be remembered that in molecular or volumetric 
equations a single molecule represents one volume of a gas 
(or vapor) and the coefficient indicates the number of vol¬ 
umes. 

EXERCISES 

1 . State and illustrate (a) Gay-Lussac’s law and (h) Avogadro’s 
law. 

2 . (a) State the argument proving that a molecule of oxygen con¬ 
sists of two atoms. (6) Apply (a) to nitrogen. 

3 . Hydrogen and nitrogen combine in the ratio of 3 to 1 to form 2 
volumes of ammonia. Show from this relation that a molecule of 
hydrogen contains at least two atoms. 


GAY-LUSSAC’S LAW 291 

4. What is the relation between molecular weight and vapor 
density ? 

6 . Why is the formula of water vapor H 2 O and not HO or H 2 O 2 ? 

6 . How are the molecular weights of gases determined by (a) the 
vapor density method, and (6) the molar method? 

7. Define and illustrate (a) molecular equation, (6) mole, (c) gram- 
molecular weight, (d) gram-molecular volume. 

8. What is a molecular formula? What is the molecular formula 
of oxygen, nitrogen, chlorine, hydrogen? 

9. How is a molecular formula determined? Illustrate. 

10. Express the following as molecular equations: (a) One volume 
of phosphorus vapor and six volumes of chlorine form four volumes of 
phosphorus trichloride (PCI3) vapor; (6) carbon and water (vapor) 
form hydrogen and carbon monoxide. 

PROBLEMS 

1 . 1,000,000 molecules of hydrogen will unite with how many mole¬ 
cules of oxygen to form how many molecules of water vapor ? What 
will be the relative weights of hydrogen and water vapor? 

2. A liter of sulfurous oxide gas (SO 2 ) weighs 2.9 gm. Calculate the 
molecular weight of this compound. Compare with the theoretical 
molecular weight. 

3. The vapor density of hydrogen chloride is 1.14. Calculate the 
molecular weight. 

4. If 3000 cc. of an oxide of carbon weigh 3.75 gm., what is the 
molecular weight, formula, and name of the oxide ? 

6 . Calculate the correct formula of the compound corresponding 
to (a) C = 85.71, H = 14.29, vapor density = 2.1875; (6) C = 39.9, 
H = 6.7, O = 53.4, vapor density = 1.906. 

6. What volume of the constituent gases can be obtained by the 
complete decomposition of 6 1. of ammonia? 

7. Write the equation for the reaction between nitric oxide (NO) 
and oxygen. What volume of oxygen is needed for 10 1. of nitric oxide ? 
What will be the volume of the product ? 

8 . The correct formula of a gas is CH4. (a) How many grams does 

a mole weigh ? (5) What volume does this weight occupy ? (c) What 

is the weight of 11. of the gas? 

9. A gas contains 69.49 per cent of oxygen and 30.51 of nitrogen. 
(a) What is its simplest formula? (6) If 500 cc. weigh 2.042 gm., what 
is the correct formula? 

SUGGESTIONS FOR LABORATORY WORK 
(References are to Newell’s Laboratory Exercises in Chemistry) 
Exercise 41 — Weight of 22.4 Liters of Oxygen. 


292 


A BRIEF COURSE IN CHEMISTRY 


TOPIC XII: FINDING ATOMIC WEIGHTS 

376. Atomic weights. — According to Dalton^s atomic 
theory (§ 356), atoms have the same weight if alike, but a 
different weight if different. And the relative weights of 
atoms are called atomic weights. 

The standard for atomic weights is oxygen = 16. 

376. Finding approximate atomic weights from molecular 
weights. — Approximate atomic weights can be determined 
from molecular weights. 

The molecular weights of compounds can be determined 
by experiment (§§ 367, 371). After the molecular weights of 
several related compounds have been determined, we can 
readily find the approximate atomic weight of the constituent 
elements. That is, it is only necessary to apportion the 
molecular weights among the respective elements of the 
compounds, and then select the minimum weight in each 
case. 

For example, let us apply this procedure to finding the 
atomic weight of carbon. The steps are as follows: 

(1) Find by experiment the molecular weights of several 
compounds containing carbon. 

(2) Find by analysis the per cent of carbon in each com¬ 
pound. 

(3) Find the portion of carbon in each molecular weight 
by multiplying the molecular weight by the corresponding 
per cent of carbon. 

(4) Select the minimum value as the approximate atomic 
weight. 

The steps are embodied in Table XIII. In the table, 
column I contains the name of each compound, column 2 the 
corresponding molecular weight, column 3 the per cent of 
carbon in each compound, and column 4 the weight of carbon 
in the corresponding molecular weight. The weights in 
column 4 are found by multiplying the molecular weight 
(in column 2) of the compound by the corresponding per 
cent of carbon (in column 3); thus 28 times 0.429 = 12. 
The minimum value 12 is the atomic weight of carbon. 

Obviously, the minimum weight must be the weight of a 


FINDING ATOMIC WEIGHTS 


293 


TABLE XIII. — Determination of the Atomic Weight 
OF Carbon 


Compound 

Mol. 

Wt. 

Per 

Cent 

OP 

Car¬ 

bon 

Wt. 

OP 

Car¬ 

bon 

Compound 

Mol. 

Wt. 

Per 

Cent 

OP 

Car¬ 

bon 

Wt. 

OP 

Car¬ 

bon 

Carbon Monoxide 

28 

42.9 

12 

Ethylene. 

28 

85.7 

24 

Carbon Dioxide. 

44 

27.3 

12 

Acetylene 

26 

92.3 

24 

Methane . . . 

16 

75.0 

12 

Ether. . 

74 

64.9 

48 

Ethane.... 

30 

80.0 

24 

Propane . 

44 

81.8 

36 



single atom, for it is highly probable that one or more com¬ 
pounds in a representative group will contain only one atom 
of a given element; and 
the part of the molecu¬ 
lar weight apportioned 
to this element will 
of course be its atomic 
weight. In the com- 
jxiunds that contain a 
multiple of this weight, 
it is likewise obvious 
that the molecule must 
contain several atoms of 
the element. Thus, the 
weight of carbon in 
ethane is twice that in 
methane, and we con¬ 
clude that a molecule of 
ethane contains two 
atoms of carbon — a con¬ 
clusion in harmony with 
other observations. 

The numerical results 
obtained in applying the 


Fig. 166. — The Italian chemist Canniz¬ 
zaro (1826-1910) whose work laid the 
foundations of our present system of 
atomic weights. 


above method to the elements oxygen, hydrogen, chlorine, 
nitrogen, and carbon are summarized in Table XIV. In 
this table, for the sake of simplicity, whole numbers are 


















294 


A BRIEF COURSE IN CHEMISTRY 


used (except in the case of chlorine) and per cents are 
omitted. The minimum weight in each case is the atomic 
weight of the element, e.g., Cl = 35.5. 

This method of determining approximate atomic weights 
was proposed about 1858 by the Italian chemist Cannizzaro 
(Fig. 166). 

TABLE XIV. — Determination of Approximate 
Atomic Weights 


Compound 

Molec¬ 

ular 

Wt. 

Wt. of 
Oxy¬ 
gen 

Wt. of 
Hydro¬ 
gen 

Wt. of 
Chlo¬ 
rine 

Wt. of 
Nitro¬ 
gen 

Wt. of 
Car¬ 
bon 

Water. 

18 

16 

2 

_ 

_ 

__ 

Hydrogen Peroxide . . 

34 

32 

2 

— 

— 

— 

Hydrogen Chloride . . 

36.5 

— 

1 

35.5 

— 

— 

Ammonia. 

17 

— 

3 

■ — 

14 

— 

Nitric Acid. 

63 

48 

1 

— 

14 

— 

Nitrous Oxide .... 

44 

16 

— 

— 

28 

— 

Nitric Oxide .... 

30 

16 

— 

— 

14 

— 

Nitrogen Dioxide . . . 

46 

32 

— 

— 

14 

— 

Carbon Monoxide. . . 

28 

16 

— 

— 

— 

12 

Carbon Dioxide . . . 

44 

32 

— 

— 

— 

12 

Methane. 

16 

— 

4 

— 

— 

12 

Ethylene. 

28 

— 

4 

— 

— 

24 

Acetylene. 

26 

— 

2 

— 

— 

24 

Ether. 

74 

16 

10 

— 

— 

48 

Ethyl Alcohol .... 

46 

16 

6 

— 

— 

24 

Chloroform. 

119.5 

— 

1 

106.5 

— 

12 

Carbon Tetrachloride 

154 

— 

— 

142 

— 

12 

Cyanogen Chloride . . 

61.5 

— 

— 

35.5 

14 

12 

Minimum weight of each 
element. 


16 

1 

35.5 

14 

12 


377. Exact atomic weights. — Atomic weights obtained 
from approximate molecular weights are, of course, approxi¬ 
mate. When the approximate atomic weight of an element 
has been chosen by the method described in § 376, its exact 
atomic weight is determined by accurate chemical analysis 
of compounds prepared from carefully purified substances. 
























FINDING ATOMIC WEIGHTS 


295 


The general method can be illustrated by an actual case kindly fur¬ 
nished by the American chemist Richards (Fig. 167), who made excep¬ 
tionally accurate determinations of atomic weights. In determining the 
atomic weight of chlorine he found that 28.26299 gm. of silver chloride 
were formed from 21.27143 gm. of silver. He accepted AgCl as the 
formula of silver chloride and 107.880 as the atomic weight of silver, and 
calculated the atomic weight of chlorine thus: — 

28.26299 - 21.27143 = 6.99156 

Wt. of silver: Wt. of chlorine:: At. wt. of silver: At. wt. of chlorine 
21.27143 ; 6.99156 :: 107.880 : x 

X = 35.458 

378. International atomic weights. — An international 
committee formerly selected the most accurate atomic weight 
of each element (except 
oxygen, of course, which 
was adopted as 16 sev¬ 
eral years ago). These 
weights were embodied 
in a table published at 
frequent intervals and 
called the International 
Table of Atomic Weights. 

The table given on the 
inside of the back cover 
of this book is not an 
international table but a 
table of the most recent 
exact weights (together 
with other data) com¬ 
piled by American chem¬ 
ists. In this table the 
accepted atomic weights 
are placed in one column 
and the corresponding 
approximate values 
(selected by the author) in another column. The approxi¬ 
mate atomic weights in the table are sufficiently accurate for 
general use, e.g., in making chemical calculations and in 
solving the problems in this book (unless otherwise directed). 





296 


A BRIEF COURSE IN CHEMISTRY 


EXERCISES 

1 . State the four steps in finding atomic weights by the “ minimum 
weight ” method. Illustrate by carbon. 

2 . Why is 16 the atomic weight of oxygen? 

3 . What is the atomic weight of (a) Al, Ba, Br, Ca, C, Cl, Cu, F, 
Au, H; (6) I, Fe, Pb, Mg, Mn, Hg, N, O, P, K; (c) Si, Ag, Na, S, Sn, 
Zn? 


PROBLEMS 

1 . If 2 gm. of potassium chloride yield 3.84 gm. of silver chloride, 
calculate the exact atomic weight of potassium (assume Cl = 35.457). 

2 . In a synthesis of hydrogen bromide, 0.8606 gm. of hydrogen com¬ 
bined with 68.25033 gm. of bromine. Calculate the exact atomic 
weight of bromine (assume H = 1.008). 


TOPIC XIII: NITROGEN OXIDES 

379. Nitrogen oxides. — The three important nitrogen 
oxides are nitrous oxide (N2O), nitric oxide (NO), and nitro¬ 
gen dioxide (NO2). There is also an oxide called nitrogen 
tetroxide (N2O4). 

380. Nitrous oxide. — This gas is prepared by gently 
heating ammonium nitrate. The equation is : — 

NH4NO3 = N2O + 2H2O 

Ammonium Nitrate Nitrous Oxide Water 

The gas is colorless, and has a faint but pleasant odor. 
It is soluble in water, and the solution has a sweet taste. 

Nitrous oxide does not burn, but it supports the com¬ 
bustion of many well-burning substances, though not so 
vigorously as oxygen does. Thus, sulfur, if burning well, 
will burn in nitrous oxide. In its power to support combus¬ 
tion it resembles oxygen (§ 27). It is distinguished from 
oxygen by its failure to form brown fumes (NO 2 ) when 
mixed with nitric oxide (§ 381). 

The most striking property of nitrous oxide is its effect 
on the human system. If inhaled for a short time, it causes 
more or less nervous excitement. If breathed in large quan¬ 
tities, it produces temporary unconsciousness and insensi- 


NITROGEN OXIDES 297 

bility to pain. The gas, mixed with a small proportion of 
air or oxygen, is often used as an anaesthetic in dentistry. 

381. Nitric oxide. — This gas is usually prepared by the 
interaction of copper and dilute nitric acid (sp. gr. 1.2). 
The equation is : — 

3 Cu + 8 HNOs = 2 NO + 3 Cu(N 03)2 + 4 H 2 O 

Copper Nitric Acid Nitric Oxide Copper Nitrate Water 

Nitric oxide is a colorless gas. It is a little heavier than 
air and only slightly soluble in water. 

Upon exposure to air {e,g., in preparing it in a test tube 
or flask), it combines at once with the oxygen, forming 
reddish brown fumes of nitrogen dioxide — a striking change. 
The equation is : — 

2 NO -h O2 = 2NO2 

Nitric Oxide Oxygen Nitrogen Dioxide 

This property distinguishes nitric oxide from all other gases. 
It does not burn, nor support combustion, unless the burning 
substance (e.gr., phosphorus or sodium) introduced is hot 
enough to decompose the gas into nitrogen and oxygen, and 
then the liberated oxygen assists the combustion. 

382. Nitrogen dioxide. — This is the reddish brown gas 
formed by the direct combination of nitric oxide and oxygen. 
It is also produced by heating certain nitrates, thus: — 

2Pb(N03)2 = 4 NO 2 + 2PbO + O 2 

Lead Nitrate Nitrogen Dioxide Lead Oxide Oxygen 

The fumes of nitrogen dioxide appear when nitric acid and 
metals interact, but, as stated above, the nitrogen dioxide 
is produced by a second reaction, viz., the combining of 
nitric oxide with the oxygen of the air. 

Nitrogen dioxide has a disagreeable odor, and it is poi¬ 
sonous if breathed in moderate quantities. It interacts with 
water and yields under ordinary conditions nitric acid and 
nitric oxide, thus : — 

3 NO 2 + H 2 O = 2 HNO 3 + NO 

Nitrogen Dioxide Waiter Nitric Acid Nitric Oxide 


298 


A BRIEF COURSE IN CHEMISTRY 


It also dissolves in concentrated nitric acid, forming fuming 
nitric acid, which is a vigorous oxidizing agent. 

When the reddish brown gas is cooled, it gradually loses 
color and at about 26° C. becomes a yellow gas. This yellow 
gas has the composition represented by the formula N 2 O 4 and 
is called nitrogen tetroxide. 

EXERCISES 

1 . How are (a) nitrous oxide, (6) nitric oxide, and (c) nitrogen dioxide 
prepared ? 

2 . State the characteristic properties of the three nitrogen oxides. 

3 . What is the test for each oxide ? 

SUGGESTIONS FOR LABORATORY WORK 

(References are to Newell’s Laboratory Exercises in Chemistry) 

Exercise *39 — Nitric Oxide and Nitrogen Dioxide. 

Exercise S36 — Nitrous Oxide — T. 


TOPIC XIV: FUELS —FLAMES 

383. Measurement of fuel value. — Besides the British 
thermal unit (B. t. u.) and the small calorie (cal.) (§ 236), 
there is another heat unit called the large calorie (Cal.). It 
is one thousand times the small calorie, and is used when the 
heat value is large. For example, the heat value, often 
called the fuel value, of different kinds of food is usually 
expressed in large calories. Thus, one gram of starch in 
burning gives out 4 Calories and one gram of fat 9 Calories 
(§ 405). 

384. The calorimeter is used to find fuel value. — The 

fuel value of coal, charcoal, and food (dried) is measured by 
burning a weighed quantity of the solid in a calorimeter. 

The essential part of the apparatus is a strong metal vessel, called a 
bomb, which is immersed in another vessel containing a known weight 
of water. The heat from the burning substance raises the temperature 
of the water and the rise is carefully measured by an accurate thermom¬ 
eter. From the weight of the substance, the weight of the water, the 


FUELS — FLAMES 299 

rise in temperature, and certain allowances for the apparatus itself, 
the fuel value of the substance can be computed. 

386. Producer gas. — This is made by forcing air (some¬ 
times together with steam) through a deep coal (or coke) fire 
in a special kind of furnace (§ 65, Fig. 168). If steam is used, 
the gaseous product contains 25 to 30 per cent of carbon 
monoxide, 50 to 60 per cent of nitrogen, and 10 to 13 per cent 
of hydrogen. Its fuel value is 
rather low, about 145 B. t. u. 
per cubic foot, though it varies 
with the process. Producer 
gas is used as a fuel in many 
industrial and metallurgical 
operations. (See also § 66 .) 

386. Water gas. — This is 
made by forcing steam through 
a thick bed (6 or 7 feet) of hot 
coke or anthracite coal. It 
contains hydrogen and carbon 
monoxide — from 40 to 50 per 
cent of each — and is used to 
some extent as fuel. More 
often it is enriched by spray¬ 
ing in petroleum oil, and then 
used as an illuminant. If 
used as fuel, enriched water gas 
has a fuel value of 500 to 600 
B. t. u. per cubic foot. 

387. Manufacture of water gas. 

— The essential parts of the appa¬ 
ratus are shown diagrammatically 
in Fig. 169. (1) Air is forced by a 

blower {A) through the fire in the 
generator {B), the hot gases pass 
down the carburetor (C), up into the superheater (D), and escape 
through an opening (not shown) into the air. This operation is called 
the “ blow.’’ It heats the fire brick inside the carburetor and super¬ 
heater intensely hot; air is often forced in to raise the temperature. 

(2) The air valves and the opening at the top of the superheater are 
now closed, and the “ run ” begins. Steam is forced into the generator 



Fig. 168. — Sketch of furnace for 
making producer gas. Coal 
enters through A, B, and air 
through D into C. Producer gas 
escapes through F and ashes are 
removed at E. 

















300 


A BRIEF COURSE IN CHEMISTRY 


at the bottom. In passing through the mass of incandescent carbon, 
the steam and carbon interact thus: — 

C + H 2 O = CO + H 2 

Carbon Steam Carbon Monoxide Hydrogen 

The mixed gases rise to the top of the carburetor, where they meet a 
spray of oil. And as the gaseous mixture passes down the carburetor 
and up the superheater, the hydrocarbons (§ 246) of the oil are trans¬ 
formed by the intense heat into gaseous hydrocarbons which do not 
liquefy when the final gas is cooled. 

(3) From the superheater the water gas passes through the purify¬ 
ing apparatus {E) into a holder. A ton of hard coal yields about 
44,000 cubic feet of enriched water gas. 



388. The illuminating gas flame. — In an ordinary il¬ 
luminating gas flame (Fig. 170) the gas issues from a slit 
in the burner tip and spreads out in a thin layer. When the 
gas is ignited, the hydrocarbons decompose and the products 
burn; the hydrogen burns to water and part of the carbon 
burns to carbon dioxide. Some of the fine particles of carbon 
do not burn at once, but are heated hot and make the flame 
luminous. Eventually all the carbon burns, if sufficient 
oxygen is supplied and the temperature is high enough. 





































































FUELS — FLAMES 


301 



Fig. 170 


An illuminat¬ 
ing gas flame. 


The flame smokes or deposits free carbon, if the air supply 
is cut off or diminished, or the temperature reduced. For 
example, if a bottle is held low down 
over the flame, smoke is given off; 
the same result is produced when the 
flame is cooled. The presence of un¬ 
burned carbon in the flame can be 
shown by putting a piece of crayon 
or a glass rod in or just over the 
flame; a deposit of soot (carbon) 
soon gathers. 

The flame is flattened to expose a 
large surface to the air so that all the 
carbon will be consumed, and thereby 
increase the lighting surface. 

There are two distinct parts, or 
zones, to this flame (Fig. 170). The lower part near the tip 
is black and consists largely of cold, non-luminous gas, while 
the upper part— “ the flame ”—is yellow-white 
and contains luminous particles of carbon. 

389. Other luminous flames. — In an oil 
lamp, e.gr., kerosene lamp, the oil is drawn up 
through the wick and volatilized by heat into 
a gas, which burns; the air supply is increased 
through small holes in the burner (below the 
flame) or in large lamps through a central vent 
also. In many lighthouses the vapor from oil 
is burned directly {i.e., not through a wick as 
formerly). In a candle, the heat from the burn¬ 
ing wick melts and volatilizes the wax, and this 
gas burns. The structure of luminous flames is 
essentially alike. 

390. The candle flame. — Examination of 
the sketch in Fig. 171 reveals four somewhat 
conical portions: — 

(1) Around the wick there is a dark cone (A), filled with 
combustible, but unburned, gases formed by the decom¬ 
position of the carbon compounds in the melted wax. 

(2) Around the lower part of the dark cone is a faint 



Fig. 171. — 
Sketch of 
the parts of 
a candle 












302 


A BRIEF COURSE IN CHEMISTRY 



Fig. 172. — Charred 
paper showing the 
hottest part of a 
candle flame. 


bluish cup-shaped part {B, B). It is the lower portion of the 
exterior cone {D, D). 

(3) Above and surrounding the dark cone is the luminous 
portion (C). It is the largest and most important part of 

_the flame. Combustion is incomplete here, 

because little oxygen can pass through the 
exterior cone. The hydrocarbons undergo 
complex changes. The most characteristic 
change is the liberation of small particles 
of carbon. This carbon, heated to incan¬ 
descence by the burning gases, makes the 
flame luminous. 

(4) The exterior cone (D, D) is almost 
invisible. Here the combustion is com¬ 
plete, because oxygen of the air changes all 
the carbon to carbon dioxide. It is the 
hottest region of the flame, for by pressing a piece of stiff 
white paper for an instant down upon the flame almost to 
the wick the paper will be charred by the hot outer portion 
of the flame (Fig. 172). 

The gaseous products of the combustion of a candle are 
water vapor and carbon dioxide. A bottle held over a 
burning candle has, at first, a deposit of moisture on the 
inside; and if the bottle is removed and calcium hydroxide 
solution is added, the presence 
of carbon dioxide is shown 
by the cloudiness of the solu¬ 
tion. 

We can readily show with 
a candle flame how the lumi¬ 
nosity of hydrocarbon flames 
is affected by temperature. 

Thus, if a coil of copper wire 
is lowered upon a candle flame, the flame smokes, loses its 
yellow color, and finally goes out; but if a coil of hot wire is 
used, the flame burns unchanged (Fig. 173). 



Fig. 173. — Effect of lowering the 
temperature of a candle flame. 


It should be noted that not all luminous flames are hydrocarbon 
flames. Thus, magnesium burns with a brilliant flame. Its luminosity 
is due to the incandescence of solid particles of magnesium oxide. 









FUELS — FLAMES 


303 


Hence the use of magnesium in fireworks, e.g., star shells. Similarly, 
the bright flame of burning phosphorus is accounted for by the incandes¬ 
cent particles of solid phosphorus pentoxide (Fig. 182). 

391. Acetylene and its flame. — Acetylene (C 2 H 2 ) is a 
gas which is prepared by the interaction of calcium carbide 
and water, thus : — 

CaCi, + 2 H 2 O = C 2 H 2 + Ca(OH )2 

Calcium Carbide Water Acetylene Calcium Hydroxide 

It can be prepared in this way for use on a small scale, as in a 
miner’s lamp or portable lantern, or on a large scale for 
lighting towns, houses, caves, and mining camps. 

Acetylene burns with a luminous, smoky flame. But 
when considerable air is mixed with the gas as the latter issues 


Fig. 174. — Acetylene burner and flame. 





from a small opening, the mixture burns with a brilliant, 
white flame, which does not smoke. 

A common form of acetylene burner is shown in Fig. 174. 
The acetylene as it escapes from the supply pipe (A) into 
the burner sucks in air through the small side holes (B). 
This mixture, upon ignition, burns as a small flat flame at 
right angles to the burner (Fig. 174, right). 

392. The oxy-acetylene flame. — A mixture of acetylene 
and the proper proportion of oxygen burns with an intensely 
hot flame; a temperature of nearly 3500° C. can be reached 
if the mixture is burned by a special blowpipe called an 
oxy-acetylene torch (Fig. 175). 

The thermochemical equation is : — 

2 C 2 H 2 + 5 O 2 = 4 CO 2 + 2 H 2 O + 607,520 cal. 

Acetyleue Oxygen Carbon Dioxide Water 




304 


A BRIEF COURSE IN CHEMISTRY 


The oxy-acetylene flame is utilized in cutting ’’ and 
welding metals (Fig. 176). (See § 35.) 



Fig. 175. — Oxy-acetylene torch and flame. 



Fig. 176. — Welding an iron grill with an oxy-acetylene flame. 


EXERCISES 

1. What is (a) a British thermal unit, (b) a small calorie, and (c) a 
large calorie ? What is the abbreviation of each ? 

2. What is producer gas? How is it made? State some of its 
advantages as a fuel. 

3. Apply Exercise 2 to water gas. 

4. Make a sketch of (a) an illuminating gas flame and (6) a candle 
flame, 








SOAP 


305 


6 . Describe the two flames from the sketch made in Exercise 4 . 

6 . What is acetylene ? Write the equation for its formation. 

7. Draw an acetylene burner. Describe the flame from the 
drawing. 

8 . For what is the oxy-acetylene flame used? 

PROBLEMS 

1. An acetylene gas plant consumes 100 cubic feet an hour. How 

much calcium carbide would be used in a month of 30 days if the 
gas is burned an average of 5 hours a day ? (1 cu. ft. = 28.3171.) 

2. Calculate the B. t. u. liberated by burning 2000 cu. ft. of producer 
gas. 

3. How many calories are liberated by burning 520 gm. of acetylene? 

SUGGESTIONS FOR LABORATORY WORK 

(References are to NeweU’s Laboratory Exercises in Chemistry) 

Exercise *50 — Illuminating Gas Flame — T. 

Exercise *51 — Candle Flame — T. 

Exercise S40 — Acetylene — T. 


TOPIC XV: SOAP 

393. Manufacture of soap. — Soap is made by boiling 
fat, or oil, with sodium hydroxide solution. This process is 
called saponification. Sodium hydroxide produces hard 
soap, which is a mixture chiefly of sodium palmitate, sodium 
stearate, and sodium oleate. 

Most soaps are made by boihng the fat, or oil, and alkali in a huge 
kettle (Fig. 177). This operation produces a thick, frothy mixture of 
soap, glycerin, and alkah. At the proper time salt is added, thereby 
causing the soap to separate and rise to the top. The liquid beneath is 
drawn off, and from it glycerin is extracted. 

Some soaps are boiled again, and then mixed (if desired) with perfume, 
coloring matter, or filhng material, such as sodium sihcate, sand, or 
borax. Floating soaps are made by forcing air into the semi-solid mass 
before coohng. Flake soaps are made by driving off the water from a 
thin layer of soap solution and scraping off the dried soap. The best 
soaps are prepared so that the finished product will be free from un¬ 
changed fat or “ free alkali,” i.e., sodium hydroxide. 

394. The cleansing action of soap. — This is ascribed to 
two causes: (1) Soap hydrolyzes, i.e., interacts, with water 



306 


A BRIEF COURSE IN CHEMISTRY 


— especially hot water — and the liberated alkali (sodium 
hydroxide) acts upon the grease and oil that is usually mixed 
with the dirt. (2) Soap causes fat and grease to form a 
colloidal suspension (§§ 96, 261, 446); the minute globules 
remain suspended in water and adsorb the dirt, and the whole 
can be readily washed off. The second cause is the more 
efficient. 

396. What is soap? — Soap is an ester, i.e., a compound 
related to alcohols and organic acids. Thus, a simple ester 



Fig. 177. — Soap kettles filled with a boiling mixture of fat and sodium 
hydroxide. 


is ethyl acetate (CH3COOC2H5) which is formed by heating 
a mixture of ethyl alcohol (C2H5OH) and acetic acid 
(CH3COOH) (with a little concentrated sulfuric acid— § 264). 
Soap is a mixture of several esters, mainly glyceryl stearate 
and glyceryl palmitate. 

Some esters occur in flowers and fruits and in many cases 
give the fragrance and flavor. Many are manufactured for use 
in perfumery, candy, beverages, and cosmetics. A common 
ester is methyl salicylate, which has the flavor of wintergreen. 

SUGGESTIONS FOR LABORATORY WORK 

(References are to NeweU’s Laboratory Exercises in Chemistry) 

Exercise S42 — Soap. 

Exercise S41 — Esters. 












CELLULOSE — DERIVATIVES — PAPER 307 


TOPIC XVI: CELLULOSE —CELLULOSE DERIVATIVES — 
PAPER 

396. Cellulose. — This substance forms the cell walls of 
plants, and is, therefore, widely distributed. Wood consists 
largely of cellulose. The fibers of cotton, flax, and hemp 
plants are nearly pure cellulose, and from them are obtained 
cotton, linen, and hemp. The best qualities of filter paper 
are nearly pure cellulose. 

Cellulose is an inactive substance. It is insoluble in most 
liquids, and when eaten as a part of vegetables and grains. 



Fig. 178. — Scene in a rayon mill. 


it passes through the body unchanged. It is soluble in a mix¬ 
ture of ammonia and copper hydroxide, and is reprecipitated 
by sulfuric acid of a special strength. If the solution is 
pressed through small holes into the acid, long shiny threads 
of rayon are obtained (Fig. 178). In another process, cotton 
or wood pulp is treated with sodium hydroxide and carbon 
disulfide; the product, called viscose, is dissolved and formed 
into threads by pressing it through dies into special solutions. 
If pressed through a slit, a sheet is obtained, called cellophane, 
which is widely used as a wrapping for fruit, candy, and food. 

397. Derivatives of cellulose. — With nitric acid cellulose 
forms cellulose nitrates (Fig. 178a.). One is gun cotton. 






308 


A BRIEF COURSE IN CHEMISTRY 


It looks like ordinary cotton, and may be spun, woven, and 
pressed into cakes. It burns quickly, if unconfined, but 
when ignited by a percussion cap or when burned in a con¬ 
fined space, gun cotton explodes violently. It is used in 
blasting to some extent and especially for torpedoes and sub¬ 
marine mines. 

A mixture of gun cotton, ether, and alcohol soon becomes a 
plastic mass, which upon being rolled and carefully dried 
forms a transparent solid called smokeless powder. When 



Courtesy Eastman Kodak Co. 

Fig. 178a. — Washing and drying cotton preparatory to changing it into 
cellulose nitrate. 


exploded, it forms carbon dioxide and monoxide, nitrogen, 
hydrogen, and water vapor — all colorless gases; hence the 
name smokeless powder. 

A solution of certain cellulose nitrates in a mixture of 
alcohol and ether is called collodion. When collodion is 
poured or brushed upon a glass plate or the skin, the solvent 
evaporates, leaving behind a thin film. It is used in pre¬ 
paring photographic films and as a coating for wounds. 

Nitrocellulose is also an essential ingredient of special 
solutions called lacquers, which are extensively used for 
finishing automobiles, airplanes, furniture, and metals. 








CELLULOSE — DERIVATIVES — PAPER 309 


A mixture of camphor and cellulose nitrates is called 
celluloid, which is widely used in making photographic films. 
Colored varieties are made into ornamental objects and 
toilet articles. Non-breakable glass for windshields, ele¬ 
vator doors, etc., is made by putting sheets of celluloid be¬ 
tween sheets of glass. 

With a derivative of acetic acid, cellulose forms cellulose 
acetate, which is a transparent, tough, waterproof solid. 
It is non-combustible, and is made into motion picture films, 
coating for air-plane wings, and waterproof wrappings. 

398. Paper. — This consists chiefly of cellulose fibers 
matted, or interlaced, together. Most paper, especially 
that used for newspapers, is made from wood. The best 
qualities of paper are made from rags. 



Fig. 179. — The long machine called a Fourdrinier machine for making 

paper. 


In making paper from wood, the wood is first reduced to a pulp by 
grinding the wood upon a revolving stone or by heating it under pressure 
with sodium hydroxide or calcium bisulfite (acid calcium sulfite). The 
pulp is carefully washed, bleached, and washed again. It goes next to 
the beater in which revolving knives separate and cut the fibers of cellu¬ 
lose into finer particles. Here the filler (clay), sizing (rosin), and color¬ 
ing matter (if desired) are added and thoroughly mixed with the pulp. 
Sometimes this mixture is further treated in a refiner. The pulp, thus 
prepared, is suspended in a large volume of water (90 to 96 per cent) and 








310 


A BRIEF COURSE IN CHEMISTRY 


the thin mixture is pumped to the Fourdrinier machine on which the 
sheet is made (Fig. 179). The pulp mixture flows on to a flne wire 
cloth which moves slowly along; the water drains off and leaves the 
fibers on the wire as a thin moist layer, which is dried and pressed by 
hot rollers into a compact sheet. 


TOPIC XVII: FOOD AND NUTRITION 

399. The functions of food. — Food has two main func¬ 
tions. First, it supplies materials needed for the growth, 
maintenance, and repair of the body. Second, it provides 
energy necessary to keep the body at the proper temperature, 
as well as to enable us to move about and perform our work. 

400. Nutrients. — The parts of the food that nourish 
the body are called nutrients. And the complex process 
by which the nutrients become of use is called nutrition. 
Nutrients are chiefly derived from three groups of organic 
compounds, viz.^ carbohydrates, fats, and proteins. 

Starch and sugar are carbohydrates; nearly 70 per cent 
of our food is some form of carbohydrate. Butter, lard, and 
oils are examples of fats. Lean meat, eggs, milk, cheese, 
peas, beans, and grains contain protein. 

Water and inorganic substances (mineral matter), though 
not nutrients, are vitally connected with nutrition. 

401. Uses of nutrients. — Carbohydrates and fats are 
compounds of carbon, hydrogen, and oxygen. By digestion 
they are changed chemically into substances which can be 
oxidized readily. Thus, they provide energy, which is 
liberated partly as heat. They are fuel foods. 

On the other hand proteins are body builders, i.e., they 
furnish material for new tissue and ultimately replace worn- 
out muscle and nerve tissue. Proteins also furnish some 
energy, because part of their carbon and hydrogen is oxidized. 
They differ from carbohydrates and fats in being composed 
partly of nitrogen. 

402. Function of water and mineral matter. — Water and 
mineral matter do not build tissue or furnish energy. Water 
makes up about 70 per cent of the weight of the body. It 
keeps the tissues soft and pliable, dissolves juices and digested 


FOOD AND NUTRITION 


311 


food, and assists in eliminating poisonous matter from the 
body. 

Although mineral matter makes up only about 4 per cent 
of the weight of the body, it is indispensable for life processes. 
The mineral matter in the body supplies the materials for the 
rigid parts of the body, e.g., bones and teeth, and furnishes 
acids, bases, salts, and organo-metallic compounds, which 
give many fluids and juices of the body vital properties. 
Hydrochloric acid of the gastric juice and the hematin of 
the red blood corpuscles are examples of necessary mineral 
matter. Mineral matter is obtained from vegetables and 
fruits and from seasoning added to food. 

403. Composition of foods. — The average composition 
(in per cent) of the edible portion of some common foods is 
shown in Fig. 180 and also in Table XV. 

TABLE XV. — Composition of Foods (in Per Cent) 


Foods 

Water 

Carbo¬ 

hydrate 

Fat 

Protein 

Mineral 

Matter 

Apples. 

84.6 

14.2 

0.5 

0.4 

0.3 

Bacon. 

20.2 

— 

64.8 

9.9 

5.1 

Beans (dried) .... 

12.6 

59.6 

1.8 

22.5 

3.5 

Beefsteak (sirloin) . . 

61.9 

— 

18.5 

18.6 

1.0 

Butter. 

11.0 

— 

85.0 

1.0 

3.0 

Cheese (cream) . . . 

34.2 

■ 2.4 

33.7 

25.9 

3.8 

Codfish (fresh) .... 

82.5 

— 

0.3 

16.3 

0.9 

Com (green) .... 

75.4 

19.7 

i.i 

3.1 

0.7 

Eggs. 

73.7 

— 

10.5 

14.8 

1.0 

Grapes. 

77.4 

19.2 

1.6 

1.3 

0.5 

Ham (smoked) . . . 

40.3 

— 

38.8 

16.1 

4.8 

Mutton (forequarter) 

52.9 

— 

30.9 

15.3 

0.9 

Oatmeal. 

7.3 

67.5 

7.2 

16.1 

1.9 

Peanuts. 

9.2 

24.4 

38.6 

25.8 

2.0 

Potatoes. 

78.3 

18.4 

0.1 

2.2 

1.0 

Rice. 

12.3 

79.0 

0.3 

8.0 

0.4 

Tomatoes . . . . . 

94.3 

3.9 

0.4 

0.9 

0.5 

Walnuts (English) . . 

2.5 

16.1 

63.4 

16.6 

1.4 


404. Food as a source of energy. — The food we eat 
undergoes complex chemical changes in the body. These 
























312 


A BRIEF COURSE IN CHEMISTRY 





PEANUT 

9.2% 

25.8 

38.6 

24.4 

2.0 


iiimniuii^i 


MILK 

87.0% 

8.0 

4.0 

5.0 

1.0 




Water Protein 



Fat Carbohydrate Ash 


Fig. 180. — Average composition of common foods. 










































































































































































































FOOD AND NUTRITION 


313 


changes take place at first in the digestive organs and con¬ 
stitute the process called digestion. The digested food, 
which is in a liquid or dissolved form, is absorbed and trans¬ 
ported to the various parts of the body, where it is built up 
into the tissue of which the various organs consist. 

The tissues are constantly undergoing oxidation or some 
related process. Oxygen for this purpose comes from the air 
we breathe into the lungs; here the oxygen combines with 
the haemoglobin of the blood and in this loosely combined 
form is distributed by the blood to the organs and muscles 
of the body. The worn-out tissue and products of digestion 
are slowly transformed by the oxygen into carbon dioxide 
and water (§ 34). 

Heat is liberated by these chemical changes. We might 
call food a fuel, because its digested products are oxidized, 
just as fuels are oxidized when they burn. The heat-produc¬ 
ing power of food is called its fuel value. And just as differ¬ 
ent fuels differ in the amount of heat liberated per pound, so 
various foods differ.in their fuel value. 

406. Fuel value of food. — The fuel value of food is found 
by burning a weighed quantity of the food in a bomb calo¬ 
rimeter and measuring the amount of heat liberated (§ 384). 
The unit used in measuring the fuel value of food is the large 
Calorie (Cal.). This is the amount of heat that will raise 
the temperature of 1 kilogram of water 1° C. 

Not all the food we eat, however, is digested nor is the 
nutritive portion completely transformed in the body. 
Hence the heat values of uneaten food obtained by the bomb 
calorimeter must be changed a little to make up for these 
losses. These corrected fuel values are called physiological 
fuel values. They are the fuel values of the part of the food 
that is actually transformed into energy in the body — the 
real fuel value of the food digested and transformed. The 
physiological fuel values are the ones usually meant when 
the term fuel value is used. 

The fuel value of food is often stated in Calories per pound, 
^.6., the number of Calories furnished by one pound of food. 
The fuel values of the foods tabulated in § 403 are shown 
in Table XVL 


314 


A BRIEF COURSE IN CHEMISTRY 


TABLE XVI. — Fuel Value of Foods 
(Calories per Pound) 


Apples . . . 290 

Bacon . . . 2840 

Beans.... 1605 

Beefsteak . . 1130 

Butter . . . 3491 

Cheese . . . 1950 


Codfish . 

. . 325 

Corn . . 

. . 470 

Eggs . . 

. . 720 

Grapes . 

. . 450 

Ham . . 

. . 1940 

Mutton 

. . 1595 


Oatmeal . . . 1860 

Peanuts . . . 2560 

Potatoes. . . 385 

Rice .... 1630 

Tomatoes . . 105 

Walnuts . . . 3285 


406. How much food do we need ? — The amount of food 
needed varies with many factors, e.g.^ weight, work, age, 
and sex. A growing boy needs more than his father, a 
football player more than one who watches the game, and a 
hard laborer more than one who has a sedentary occupation. 
An average adult man who does moderate muscular work 
needs enough food to yield from 2500 to 3000 Calories a day. 
A satisfactory division among the nutrients would be: — 


Carbohydrate. 400^500 gm., or 1600-2000 Cal. 

Fat.70^85 gm., or 630-765 Cal. 

Protein.75-'§0 gm., or 300-320 Cal. 


An average adult woman needs from 2000 to 2500 Calories 
a day. 

The fuel value of food is only part of the story. Sugar 
or butter would furnish enough heat, but a diet consisting 
largely of carbohydrates or fats would not repair worn-out 
tissue. We must eat food containing protein, as well as 
carbohydrates and fat. Authorities agree that an adult who 
does moderate muscular work needs 75-80 grams (3 ounces) 
of protein per day. If more is eaten, it is partly consumed 
as fuel and partly rejected by the body. 

The amount of food we need, stated in Calories, can be 
estimated from Table XVII, which shows the approximate 
expenditure of energy (in Calories) per hour. This table 
itemizes the energy (in Calories) given out, so to speak, by a 
healthy man or woman weighing 154 lb. and 123 lb. respec¬ 
tively. By adding the Calories for the different items we 














FOOD AND NUTRITION 


315 


can get the approximate value of the food we must eat per 
day to give this energy. 


TABLE XVII. — Approximate Expenditure of Energy 
IN Calories per Hour 


Sleeping. 

. 60-70 

Awake, lying still . . 

. 70-85 

Sitting, at rest . . . 

100 

Standing, at rest . . 

115 

Typing, rapidly . . . 

140 


Walking, slowly .... 200 
Moderate exercise .... 240 

Walking, actively .... 300 

Hard exercise.480 

Running.500 


407. Vitamins. — Vitamins are present in food in minute 
amounts yet they are essential to growth and good health. 
Our diet should always include certain foods which furnish 
vitamins, especially milk, leaf vegetables, unpolished cereals, 
and fruit juices. There are several types of vitamins, but 
we can consider only the following: — 

(1) Vitamin A, or fat-soluble A, occurs in butter, milk, 
egg yolk, cod liver oil, and spinach. It is not found in meats 
(to any extent), nuts, sugars, or oils from plants. Vitamin A 
contributes to general vigor and health. It is especially 
needed to promote the growth of children. 

(2) Vitamin B, or water-soluble B, occurs in yeast, nuts, 
fruits, spinach, tomatoes, and the outer coating of rice and 
other grains. It is not found in refined flour. Vitamin B 
is essential to normal nutrition. Its lack leads to stunted 
growth. 

(3) Vitamin C, or water-soluble C, occurs in lemons, 
oranges, tomatoes, cabbage, lettuce. Nuts, meats, and 
grains (especially refined ones) contain little or none. Vita¬ 
min C prevents and cures scurvy. 

408. How to select an adequate diet. — By means of 
Table XVIII it is easy to arrange several sets of daily menus 
which will furnish sufficient fuel value and include the ap¬ 
proximate amount of protein. By learning the fuel value 
and protein content of the important items, or those used 
frequently, we can readily tell if the diet is adequate. 














316 


A BRIEF COURSE IN CHEMISTRY 


TABLE XVIII. — Servings of Food 


FOOD 

Weight in Grams 

Total 

Calo¬ 

ries 

Average 

Serving 

Protein 

Fat 

Carbo¬ 

hydrate 

Apple, baked . 

135 

1 

1 

26 

117 

Apple, raw. 

150 

0.5 

0.5 

16 

70 

Banana, raw. 

194 

1.5 

0.8 

28 

125 

Beans, baked. 

185 

19 

9 

47 

345 

Beets (4 heaping Tbs.) . . 

200 

5 

— 

15 

80 

Biscuit (3 baking powder) 

90 

8 

11 

45 

311 

Bread, white (2 slices) . . . 

76 

6.5 

1.5 

40 

200 

Butter . 

15 

0.2 

13 

— 

118 

Cheese (1 cu. in.) .... 

20 

6 

7 

— 

87 

Chicken, creamed (on toast) 

125 

17 

13 

22 

273 

Cocoa, 1 cup . 

230 

5 

6 

12 

122 

Codfish cakes (2) .... 

132 

19 

3 

17 

171 

Corn chowder (1.5 cups) . . 

333 

10 

10 

39 

286 

Cream of wheat. 

250 

5 

1 

33 

161 

Cream toast (1 slice) . . . 

148 

11 

11 

49 

339 

Custard (2 heaping Tbs.). . 

134 

7 

8 

19 

176 

Doughnut (1) ..... 

37 

2.5 

8 

20 

162 

Dressing, French (2 ts.) . . 

9 

— 

7 

— 

63 

Egg (1). 

50 

6.5 

6 

— 

80 

Figs (5). 

100 

4 

0.3 

73 

309 

Fish chowder (1 cup) . . . 

284 

38 

10 

30 

362 

Fish, stuffed and baked . . 

152 

33 

4 

7 

196 

Hominy (1 cup) . 

245 

3 

— 

32 

140 

Ice cream (2 heaping Tbs.) . 

100 

5 

10 

18 

182 

Lamb, roast (1 slice) . . . 

90 

20 

27 

— 

323 

Lettuce, 4 leaves .... 

50 

1 

— 

1 

8 

Macaroni and cheese (1 cup) 

200 

10 

14 

32 

294 

Milk and sugar for cereal . . 

73 

2 

2 

15 

86 

Milk, glass . 

220 

7 

9 

11 

153 

Onions, boiled. 

85 

2 

— 

10 

48 

Orange . 

250 

1.5 

0.3 

21 

93 

Pie, apple (i). 

126 

4 

12 

54 

340 

Potato, baked. 

130 

4 

— 

32 

144 

Potato, boiled. 

150 

4 

— 

31 

140 

Potato, sweet. 

100 

3 

2 

42 

200 

Prunes (5). 

50 

1 

— 

63 

256 

Pudding, chocolate farina 

141 

7 

7 

33 

223 

Pudding, cottage .... 

144 

6 

10 

77 

422 
































FOOD AND NUTRITION 


317 


TABLE XVIII. — Continued 


FOOD 

Weight in Grams 

Total 

Calo¬ 

ries 

Average 

Serving 

Protein 

Fat 

Carbo¬ 

hydrate 

Pudding, rice. 

133 

6 

7 

34 

223 

Rice, steamed (1 cup) . . . 

200 

4 

— 

40 

176 

Salad, fruit. 

262 

2 

1 

46 

201 

Salad, potato. 

150 

4 

— 

31 

140 

Salad, vegetable. 

157 

3 

1 

16 

85 

Sandwich, chicken .... 

70 

9 

4 

22 

160 

Sandwich, ham. 

70 

7 

10 

27 

226 

Shredded wheat (1) . . . 

29 

3 

0.5 

23 

108 

Snow pudding (2 heaping Tbs.) 

80 

4.5 

— 

12 

66 

Soup, cream of celery (1 cup) 

125 

3 

9 

5 

113 

Soup, cream of tomato (1 cup) 

125 

3 

9.5 

6.5 

123 

Spinach (1 heaping Tbs.) . . 

100 

2 

4 

3 

56 

Steak, Hamburg (1 cake). . 

50 

16 

8 

— 

136 


409. How to select a proper diet. — The carbohydrates 
and fats must be in the right quantity for fuel value and the 
protein in the proper proportion for tissue building (§ 406). 
Moreover, we ought to select the three classes of nutrients 
from a wide variety of foods in order to provide a mixed 
diet, obtain the right kind of protein, and secure indispen¬ 
sable vitamins and mineral matter. We should drink water 
freely — at least six glasses a day. We should eat regularly 
an abundance of vegetables and cereals, which yield mineral 
matter and furnish bulky ingredients, e.gr., cellulose, needed 
to assist the elimination of solid waste matter. Certain 
foods should always form a part of our diet, e.g., leaf vege¬ 
tables (especially spinach), milk, whole wheat, unpolished 
rice, and fruit because they provide vitamins. 

EXERCISES 

1 . Define the terms food, nutrition, and nutrients. 

2 . State the composition and fuel value of (a) bread, (6) butter, 
(c) potatoes, (d) eggs, (e) milk. 


















318 


A BRIEF COURSE IN CHEMISTRY 


3 . Select from the table in § 403 foods rich in (a) protein, (6) carbo¬ 
hydrate, and (c) fat. 

4 . Assuming that one square inch represents 1000 Calories, draw 
diagrams of the fuel value of five foods from the table in § 406 . 

6 . Use the table in § 408 for the following: (a) Make out a menu 
for breakfast, dinner, and supper which contains 75-80 gm. of protein 
and furnishes 2500-3000 cal. (b) As in (a) for 100 gm. of protein and 
about 3500 Cal. (c) As in (a) for 100 gm. of fat, 85 of protein, and 
about 4000 Cal. 

6 . Use the table in § 408 to keep a record (or make an estimate) of 
the approximate amount of carbohydrate, fat, and protein you eat in 

(a) one day, (b) three days, (c) one week, (d) one lunch. 

7. Calculate the weight of (a) mineral matter in 1 lb. of cheese, 

(b) water in 1 lb. of potatoes, (c) fat in 1 lb. of beefsteak, (d) protein in 
1 lb. of beans, (e) water in 1 lb. of butter, (/) carbohydrate in 1 lb. of 
oatmeal. 

8. (a) What weight of bacon is needed to furnish as much fat as 1 lb. 
of butter? (b) Of tomatoes to provide the carbohydrate in 1 lb. of 
potatoes? (c) Of oatmeal to equal the protein in 1 lb. of beans? 

SUGGESTIONS FOR LABORATORY WORK 

(References are to Newell’s Laboratory Exercises in Chemistry) 

Exercise S43 — Nutrients in Food — T. 

Exercise 56 — Testing Baking Powders. 


TOPIC XVIII: PHOSPHORUS —ARSENIC INSECTICIDES — 
ALLOYS OF ANTIMONY AND BISMUTH 

410. Occurrence of phosphorus. — Free phosphorus is not 
found in nature, but phosphates are numerous and abun¬ 
dant. The most common phosphate is calcium phosphate 
(Ca 3 (P 04 ) 2 ), which is the chief ingredient of “ phosphate 
rock and bones (80 per cent). Small amounts of phosphates 
occur in fertile soils and some iron ores. Complex organic 
phosphorus compounds are found in the germs of seeds and 
the nerves, brains, and muscles of animals. 

411. Manufacture of phosphorus. — Phosphorus is manu¬ 
factured in an electric furnace (Fig. 181). 

The mixture of calcium phosphate, carbon, and sand is introduced at 
A and fed continuously into the furnace by the screw B. An electric 
current passed between the electrodes EE produces the intense heat 


PHOSPHORUS 


319 


needed for the chemical change. The phosphorus vapor escapes through 
C into a condenser; the liquid residue, which is essentially calcium sili¬ 
cate, is drawn off as slag at D. The equation is; — 

2 Ca 3 (P 04)2 + 6 SiOa -f- 10 C = P 4 -f- 10 CO d- 6 CaSiOs 

Calcium Sand Carbon Phosphorus Carbon Calcium 

Phosphate Monoxide Silicate 

The phosphorus is purified, melted, and cast into small sticks. 

412. Properties of white phosphorus. — There are two 
kinds of phosphorus — commonly called white and red. 
The phosphorus prepared by the 
method just described is the 
white or ordinary form. It is a 
slightly yellow, translucent solid. 

The color deepens by exposure 
to light. At ordinary tempera¬ 
tures it is soft like wax, and can 
be cut with a knife. At low 
temperatures it is brittle. It dis¬ 
solves in carbon disulfide. 

White phosphorus is a danger¬ 
ous substance. When exposed to 
air it oxidizes quickly and gives 
off white fumes. At about 35° C. 
it takes fire, often suddenly, and 
burns with a brilliant flame. It 
must be kept beneath water and should not be touched with 
the hands unless it is wet. The safest way is not to touch it 
at all, but to use wet forceps in holding or transferring it. 
Unusual care should be taken not to leave bits of phos¬ 
phorus in deflagrating spoons or dishes. 

In moist air white phosphorus is slightly luminous. This 
property gave the element its name (from a Greek word 
meaning “ light bringer ”). 

White phosphorus has a faint odor. The fumes are very 
poisonous, and cause a dreadful disease, which rots the bones, 
especially the jaw bones (§ 414). 

White phosphorus in burning forms a cloud of white 
particles of phosphorus pentoxide (P2O5), which soon settles. 
In rather moist air, however, the pentoxide unites with 


A 



of phosphorus from calcium 
phosphate. 















320 


A BRIEF COURSE IN CHEMISTRY 


water to form phosphoric acid, which remains suspended 
as a fog. In the World War phosphorus was used to make a 
smoke screen and as an ingredient of incendiary bombs 
(Fig. 182). 

413. Red phosphorus. — This is a red solid. - It is made 
by heating white phosphorus to 250°-300° C. in a closed 
vessel freed from air. Red phosphorus is quite different 
from white phosphorus. It does not glow in the air, nor 



Fig. 182. — A phosphorus bomb dropped from an airplane and exploding 
on a warship. 

does it ignite until heated to about 240° C. It is not poison¬ 
ous, and does not dissolve in carbon disulfide. It does not 
combine with oxygen at ordinary temperatures and can be 
handled safely. 

414. Matches. — Phosphorus was formerly used ex¬ 
tensively in the manufacture of matches. Phosphorus 
sulfide (P 4 S 3 ) is now used in the United States as a substitute 
for the element. This change was made on account of a 
prohibitive tax upon phosphorus matches (two cents per 
hundred matches); the tax was levied mainly to protect 
the workmen from the disease caused by breathing phos- 





PHOSPHORUS 


321 


phorus fumes. The sulfide is not poisonous and is as suitable 
for match heads as phosphorus itself. 

Ordinary matches are made by dipping one end of the match sticks 
first into melted paraffin and then into the “ phosphorus mixture.” 
The latter consists usually of different proportions of phosphorus sulfide, 
manganese dioxide or another oxidizing substance, and glue or some 
other binding material mixed •with a little coloring matter. These 
matches are the ordinary friction kind. By rubbing them on a rough 
surface the friction generates enough heat to ignite the phosphorus 
compound, which continues to burn owing to the oxygen supplied 
(mainly) by the oxidizing agent. The heat thereby produced sets fire 
to the paraffin, and this in turn kindles the wood. 

In safety matches the head is usually a colored mixture of antimony 
sulfide, potassium chlorate, and glue, while the surface on the box 
upon which the match must be rubbed to ignite is a mixture of red phos¬ 
phorus, glue, and powdered glass. 

416. Uses of arsenic compounds. — Arsenic trioxide 
(AS 2 O 3 ) is used to a limited extent in making fly and rat 
poison, glass (especially plate glass), and arsenic compounds 
(e.g., insecticides), and for destroying weeds and preserving 
skins in museums. 

Arsenic trioxide is a poison. The antidote for arsenic 
poisoning is fresh ferric hydroxide, which is made by adding 
ammonium hydroxide to a ferric salt, e.g., ferric chloride. 
The ferric hydroxide forms an insoluble substance with 
the arsenic compound. 

The chief use of other arsenic compounds is for de¬ 
stroying insects. The insecticides, as they are often called, 
are used in the form of dust or sprays. Paris green 
(Cu 3 (As 03)2 . Cu(€ 211302 ) 2 ) and lead arsenate (Pb 3 (As 04 ) 2 ) 
are used extensively to exterminate potato bugs and other 
insect pests. Calcium <:rsenate (Ca 3 (As 04 ) 2 ) is used to 
exterminate the boll weevil, an insect which destroys the 
cotton plant. It is dusted on from an airplane (Fig. 183). 

416. Alloys of antimony. — Antimony and bismuth are 
used to make alloys. When antimony is melted with some 
metals, especially lead and tin, the metals dissolve one an¬ 
other. Such a metallic solution upon solidifying forms 
an alloy. Alloys have different, often very different, prop¬ 
erties from the original metals. Thus, an alloy of 10 per cent 


322 


A BRIEF COURSE IN CHEMISTRY 



Fig. 183. — Airplane dusting a cotton field with calcium arsenate. 


antimony, 70 per cent lead, and 20 per cent tin expands on 
cooling, and is harder than lead. This alloy is used as type 

metal, because it makes 
the face of the type hard 
and reproduces sharply 
the dots and fine lines 
(of the mold). Babbitt 
metal contains antimony, 
tin, and copper. It is 
used for the bearings of 
machines to reduce fric¬ 
tion. Other anti-friction 
alloys contain antimony, 
lead, tin, and a little cop¬ 
per. Another alloy con¬ 
taining an appreciable 
proportion of antimony 
is Britannia metal, which 
is used in making orna¬ 
ments and table ware. 

Fig. 184. — Sprinkler head, fusible link, 417. AlloyS of bismuth. 

and fireproof door (held in place by a Biqrniith impd in 

link of fusible metal). Fusible metal is -DlSmUtn IS USeU in 

at A. making alloys which have 

























ALLOYS OF BISMUTH 


323 


low melting points. The metal itself melts at 270° C. But 
alloys of bismuth, lead, and tin melt at a much lower tem¬ 
perature. For example, Newton’s metal melts at 94.5° C. 
and Rose’s metal at 98.3° C.; while Wood’s metal, which 
contains the metal cadmium also, melts at only 60.5° C. 

These alloys of bismuth are called fusible metals. They 
are used in making safety plugs for steam boilers, fuses for 
electrical appa¬ 
ratus, and as 
connecting 
links to hold in 
place automatic 
fireproof doors 
and to close the 
heads in the au¬ 
tomatic sprink¬ 
ling apparatus 
frequently in¬ 
stalled in large 
buildings (Figs. 

184, 185 — fusible metal at A). In case of fire, the heat 
soon melts the fusible metal in the sprinkler heads, thereby 
providing openings in the pipes from which the water flows. 
The fireproof door is kept open by a weight until the heat 
melts the fusible metal and lets the door close automatically. 
(Study Fig. 184.) 

EXERCISES 

1 . Describe the manufacture of phosphorus. 

2 . State the properties of (a) white and (b) red phosphorus. 

3 . State the uses of alloys of (a) antimony and (&) bismuth. 

PROBLEMS 

1 . Calculate the weight of phosphorus in 40 tons of calcium phos¬ 
phate. 

2 . Calculate the percentage composition of (a) arsenic trioxide and 
arsenic pentoxide, and (6) antimony trichloride and antimony penta- 
chloride. 

3 . How many gm. of phosphorus can be made by the electrothermal 
process from a ton of calcium phosphate (70 per cent pure) ? 



(left). The parts of the dismantled head are shown 
(left). Fusible metal in place at A (right). 
































324 


A BRIEF COURSE IN CHEMISTRY 


TOPIC XIX: ARRANGEMENT OF THE ELEMENTS BY 
ATOMIC WEIGHTS AND BY ATOMIC NUMBERS 

418. Arrangement of the elements. — Elements differ 
from one another in many ways. Some of them, however, 

are similar in their prop¬ 
erties and can be put into 
the same class. Thus, 
there are the large classes 
of metals and non-metals. 
Certain elements are so 
closely related they can 
be put into small groups, 
e.g., fluorine, chlorine, bro¬ 
mine, and iodine (§ 440). 
Other elements can be 
similarly grouped accord¬ 
ing to their properties, 
e.g., calcium, strontium, 
and barium. 

419. Arrangement of 
the elements according 
to their atomic weights. 
— The Russian chemist 
Mendelejeff (Fig. 186) 
was the first to show how 
to arrange the elements according to a definite plan. His 
arrangement, which appeared in 1869, is based on the relation 
between the properties of the elements and their atomic 
weights. This arrangement was made into a table known 
as the periodic table of the elements. It guided chemists 
in their work for over half a century. 

Mendelejeff’s scheme of arrangement is substantially as 
follows : If the elements — omitting hydrogen and beginning 
with helium — are arranged in the order of their increasing 
atomic weights, a series results in which similar or closely 
related elements occur at definite intervals. Thus, we would 
have the accompanying arrangement of pairs for the first 
sixteen elements: 



Fig. 186. — The Russian chemist Mende¬ 
lejeff, who made the first systematic ar¬ 
rangement of the elements. 



ARRANGEMENT OF ELEMENTS 


325 


Helium 4 Lithium 7 Beryllium 9 Boron 11 

Neon 20 Sodium 23 Magnesium 24 Aluminum 27 


Carbon 12 Nitrogen 14 Oxygen 16 Fluorine 19 

Silicon 28 Phosphorus 31 Sulfur 32 Chlorine 35.5 


Examination shows, for example, that the two similar ele¬ 
ments fluorine and chlorine occur as a vertical group, fluorine 
being the eighth and chlorine the sixteenth. The elements 
in the other groups, especially lithium and sodium, nitrogen 
and phosphorus, are similarly related. 

An analogous relation exists among the 80 or more elements 
(with three exceptions), though not so typically as among 
the flrst twenty elements. The recurrence of similar ele¬ 
ments at definite intervals means that the long series made 
up of all the elements breaks up into shorter series of eight, 
or some number near eight. These shorter series are called 
periods, and the arrangement of the elements in this way is 
called the periodic arrangement. 

The fundamental relation of this arrangement, viz., re¬ 
currence of elements at definite intervals, is sometimes stated 
thus: — 

The 'properties of the elements var'y periodically with their 
atomic weights. 

If the long series is divided into periods and the periods 
are placed below each other, a table called the periodic table 
is obtained which is substantially like Table XIX. The 
horizontal rows numbered 1 to 6 are the periods. The ver¬ 
tical columns numbered 0 to VIII are the groups. 

420. Groups and families. — Elements which resemble 
one another are found in the same group. In some groups 
certain elements are more closely related than others, giv¬ 
ing rise to sub-groups or families. Groups and families 
are found in Table XIX, and the important ones should 
be learned. 

421. The old periodic law. — Not all the periods are so 
typical as 1 and 2, nor are the elements in some groups so 
closely related as in group I. Nevertheless a careful and 
comprehensive study of all the elements shows that in many 


TABLE XIX. — Arrangement of the Elements by Atomic Weights 


326 


A BRIEF COURSE IN CHEMISTRY 


VIII 




Iron 

Fe, 55.84 

Cobalt 

Co, 58.94 

Nickel 

Ni, 58.69 

Ruthenium 

Ru, 101.7 

Rhodium 

Rh, 102.9 

Palladium 

Pd, 106.7 

Osmium 

Os, 190.8 

Iridium 

Ir, 193.1 

Platinum 

Pt, 195.2 


VII 

A B 

Fluorine 
F, 19.0 

Chlorine 
Cl, 35.457 

Manganese 

Mn, 54.93 

Bromine 

Br, 79.92 

Iodine 

I, 126.93 




VI 

A B 

1 

Oxygen 
0, 16.0 

Sulphur 
S, 32.06 

Chromium 

Cr, 52.01 

Selenium 
Se, 79.2 

Molybdenum 

Mo, 96.0 

Tellurium 

Te, 127.5 

Tungsten 

W, 184.0 


Uranium 

U, 238.14 

> 

A B 

Nitrogen 
N, 14.0 

Phosphorus 
P, 31.027 

Vanadium 

V, 50.96 

Arsenic 
As, 74.96 

Columbium 
Cb, 93.1 

Antimony 
Sb, 121.76 

Tantalum 

Ta, 181.5 

Bismuth 

Bi, 208.0 



IV 

A B 

Carbon 
C, 12.0 

Silicon 
Si, 28.06 

Titanium 

Ti, 47.90 

Germanium 
Ge, 72.60 

Zirconium 

Zr, 91.22 

Tin 
Sn, 118.7 

Cerium* 

Ce, 140.13 

Lead 
Pb, 207.2 

Thorium 

Th, 232.12 

l-H 

A B 

Boron 

B, 10.82 

Aluminum 
Al, 26.97 

Scandium 

Sc, 45.1 

Gallium 
Ga, 69.72 

Yttrium 

Yt, 88.93 

Indium 
In, 114.8 

Lanthanum 
La, 138.9 

Thallium 
Tl. 204.0 



HH 

A B 

Beryllium 

Be, 9.02 

Magnesium 
Mg, 24.32 

Calcium 

Ca, 40.07 

Zinc 
Zn, 65.38 

Strontium 

Sr, 87.63 

Cadmium 
Cd, 112.4 

Barium 

Ba, 137.36 

Mercury 
Hg, 200.6 

Radium 

Ra, 225.97 

HH 

A B 

Lithium 

Li, 6.94 

Sodium 

Na, 22.997 

Potassium 
K, 39.10 

Copper 
Cu, 63.57 

Rubidium 
Rb, 85.44 

Silver 
Ag, 107.88 

Caesium 

Cs, 132.81 

Gold 
Au, 197.2 



o 


Helium 
He, 4.0 

Neon 

Ne, 20.183 

Argon 

A, 39.94 

Krypton 
Kr, 82.9 

Xenon 

Xe, 130.2 

Radon 

Rn. 222 

Group 

Family 

Period 1 

Period 2 

Period 3 

Period 4 

Period 5 

Period 6 


* Between Cerium and Tantalum a number of elements are omitted 






































































































ARRANGEMENT OP ELEMENTS 


327 


cases their properties vary periodically with the atomic 
weight. Mendelejeff summarized these facts in a form 
sometimes called the old periodic law, thus: — 

The properties of the elements are periodic functions of their 
atomic weights. 

The term function as used here means the exhibition of a 
special relation, viz., that of properties to atomic weight. 
Strictly speaking, the relation summarized by Mendelejeff 
is not sufficiently accurate to be called a law. Interpreted 
freely, the facts at the basis of Mendelejeff’s periodic arrange¬ 
ment mean : (1) properties and atomic weights are related ; 
and (2) this relation is exhibited in many instances at regular 
intervals. 

422. Imperfections in the old periodic arrangement. — 

(1) There are vacant places (indicated by dashes). These 
probably correspond to elements not yet discovered. Three 
vacant places in the original table were filled many years ago, 
and recently other places were filled. (2) Three elements 
are out of place, e.g., argon undoubtedly belongs under neon, 
yet its atomic weight (39.94) would necessitate its exchange 
with potassium (39.10) — an absurdity! There are two other 
pairs of elements not located where their atomic weights de¬ 
mand. (3) Hydrogen lacks an acceptable place; it really 
has no place ! (4) The table lays too much stress on a single 

valence which is not always the common valence. 

There is undoubtedly a more fundamental basis for ar¬ 
ranging the elements, viz., their atomic numbers. 

423. Arrangement of the elements according to their 
atomic numbers. — The present basis of arrangement of the 
elements is their atomic numbers. 

Each element has a serial number called its atomic number. 
Atomic numbers start with hydrogen = 1 and end with 
uranium = 92. In Table XX the elements are arranged 
in the same general way as in the old periodic arrangement, 
but this time in the strict order of their atomic numbers. 
Several points should be noted. 

(1) The general arrangement into groups and periods is 
unchanged. 


TABLE XX. — Arrangement of the Elements by Atomic Numbers 


328 


A BRIEF COURSE IN CHEMISTRY 







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I, 71 Lutecium, 72 Hafnium (Celtium). 




























ARRANGEMENT OF ELEMENTS 


329 


(2) All the elements are in their proper places. That is, 
the three pairs of elements, anomalously located by their 
atomic weights, are now in the correct order. Specifically, 
(a) argon, atomic number 18, precedes potassium, atomic 
number 19, (h) cobalt (27) precedes nickel (28), and (c) tel¬ 
lurium (52) precedes iodine (53). 

(3) The rare earth elements (see footnote 2 in Table XX) 
which have atomic numbers 59 to 72 are omitted from the 
table to avoid confusion. 

(4) Hydrogen (atomic number =1) has no acceptable 
place. 

424. Determination of the atomic number of an element.— 

The atomic numbers of the elements are assigned from the 
relative value 



of their X-ray 
spectra. 

Light and X- 
rays are similar. 

Both affect a 
photographic 
plate. The es¬ 
sential difference 
is in the wave 
lengths. X-ray 
waves being very 
much shorter. 

Visible light in 
passing through 
a glass prism is 
spread out into a band of colors called a spectrum. A spec¬ 
trum is also produced by letting light fall upon a diffraction 
grating, i.e., a plane surface of glass or metal on which is ruled 
an enormous number of fine parallel lines, the lines being sep¬ 
arated by distances of the order of the wave length of light. 

Diffraction gratings cannot be ruled fine enough for X-ray 
spectra. But crystals can be used to produce X-ray spectra, 
because the atoms in crystals are arranged in successive 
rows and layers so close together that they actually form a 
grating of the fineness adapted to X-rays. 



Fig. 187. 


Sketch of the apparatus for finding 
X-ray spectra. 





















330 


A BRIEF COURSE IN CHEMISTRY 


When a high-voltage electric current is passed through a 
vacuum tube A (Fig. 187), rays called cathode rays (t.e., 
a stream of negatively charged particles of electricity called 
electrons) are given off at the cathode. If the cathode 
rays strike a solid, metallic target B (called the anti¬ 
cathode), part of their energy is transformed into X-rays, 
which emerge from the tube. These X-rays, in passing 
through a slit and falling upon the crystal C, produce the 

X-ray spectra (alluded 
to above), which may be 
photographed on a plate 
mounted beyond D. A 
given element yields a 
characteristic X-ray 
spectrum. 

The facts, substantially 
as just stated, were dis¬ 
covered and first utilized 
by the gifted young Eng¬ 
lish physicist Moseley 
(Fig. 188). 

Moseley found by ex¬ 
periment that the X-ray 
spectra of different metals 
have conspicuous lines. 
And he discovered that 
there is a regular change 
in the position of the 
lines as one metal after 
another is used in the 
order of their increasing atomic weights. By assigning a 
number to a given element, e.g., aluminum = 13, he found 
he could arrange the elements he examined in a series which 
increased numerically by 1 as he passed up the series. 
Furthermore, he found that the order is the same as that 
required in the old periodic classification — no exceptions in 
the new order. 

Moseley’s work has been extended, and now each element 
is assigned a number, according to its position in the series, 



Fig. 188. — The English physicist Moseley 
(1889-1915), whose studies of X-ray 
spectra led to the atomic numbers of 
the elements. 



ARRANGEMENT OF ELEMENTS 331 

thereby giving a succession of whole numbers from 1 (for 
hydrogen) to 92 (for uranium). These numbers are the 
atomic numbers. They are numerically equal to the num¬ 
ber of electrons surrounding the nucleus of the atom. 

426. The new periodic table and law. — The new table 
(XX) of the elements arranged according to their atomic 
numbers is almost exactly like the old one based on atomic 
weights. The only apparent difference is that all the 
elements are now in their correct periodic places, that is, 
in the correct groups, e.g., argon in group 0 and potassium 
in group I. 

The atomic number of an element is more fundamental 
than its atomic weight. Many experiments have convinced 
chemists that atomic numbers are a deep and comprehensive 
expression of the properties of elements. Hence the rela¬ 
tions expressed in the new periodic table can be summarized 
in the new periodic law thus: — 

The properties of the elements are a periodic function of 
their atomic numbers. 


EXERCISES 

1 . What is meant by (a) period, (6) group, (c) family? 

2 . Illustrate the periodic arrangement by periods 1 and 2. 

3 . Commit to memory the names of the elements in Groups 0, I, 
II, VII. 

4 . As in Exercise 2 by the proper elements in groups I and VII. 

6 . What is the atomic number of an element ? 

6 . What is the atomic number of hydrogen, oxygen, carbon, nitro¬ 
gen, chlorine, iron, calcium, sodium? 


PROBLEMS (Review) 

1 . Calculate the per cent of oxygen in (a) water, (h) potassium chlo¬ 
rate, (c) nitric acid, (d) hme. 

2 . A candle in burning forms 13.21 gm. of carbon dioxide and 5.58 gm. 
of water. How much weight did the candle lose? What volume of 
oxygen at 0° C. and 760 mm. was required? 

3 . What volume of air (free fyom carbon dioxide and water vapor) 
contains 1 gm. of nitrogen ? 

4. What weight of sulphur is contained in 500 cc. of SO 2 ? 


332 


A BRIEF COURSE IN CHEMISTRY 


6. Suppose 50 1. of nitrous oxide are decomposed into nitrogen and 
oxygen. How many volumes of the products are formed? 

6. A compound has the composition C = 10.04, H = 0.836, Cl = 
89.13, and the vapor density is 3.735. What is the molecular formula? 



TOPIC XX: FLUORINE —BROMINE —IODINE 


426. The halogen family. — Fluorine, bromine, and iodine, 
together with chlorine (§§ 120-137), constitute a typical 

family of related ele¬ 
ments often called the 
halogens. The elements 
and their analogous 
compounds have similar 
properties, differing 
mainly in degree (§ 440). 


FLUORINE 


427. Occurrence of 
fluorine. — Fluorine oc¬ 
curs only in combination 
with other elements, e.g., 
with calcium as calcium 
fluoride (fluor spar, fluo¬ 
rite, CaF 2 ). Other native 
compounds are cryolite 
(NasAlFe) and apatite 
(Ca5F(P04)3). 

428. Preparation of 
fluorine. — Fluorine was 

first isolated in 1886 by the French chemist Moissan (Fig. 
189) by the electrolysis of a mixture of anhydrous hydro¬ 
fluoric acid (H 2 F 2 ) and acid potassium fluoride (KHF 2 ). 
The experiment was difficult and dangerous owing to the 
corrosive properties of both acid and element. 


Fig. 189. — The French chemist Moissan 
(1852-1907), who first prepared the 
corrosive element fluorine. 


Fluorine can be prepared on a commercial scale by the electrolysis 
of molten acid potassium fluoride (KHF 2 ) in a graphite cell, which serves 
as the cathode. The anode is a graphite rod. It passes down through 



FLUORINE — BROMINE — IODINE 333 

a graphite cylinder (in which the fluorine collects) which dips just below 
the surface of the molten fluoride. 

429. Properties of fluorine. — Fluorine is a greenish 
yellow gas. It smells like chlorine. Chemically, fluorine is 
intensely active. It combines with most elements readily, 
the combining being accompanied by much heat and light. 
The compounds formed are fluorides. It does not combine 
with oxygen or nitrogen, while some metals, e.g., gold, plati¬ 
num, and copper, are not readily (or only slightly) attacked 
by it. Water is decomposed by it at ordinary temperature. 

430. Hydrogen fluoride. — This gas is prepared by the 
interaction of concentrated sulfuric acid and calcium fluoride. 
The experiment is performed in a lead dish because hydrogen 
fluoride is so corrosive. The equation is : — 

CaF 2 + H2SO4 = H2F2 + CaS 04 

Calcium Fluoride Sulfuric Acid Hydrogen Fluoride Calcium Sulfate 

Hydrogen fluoride forms fumes in moist air and dissolves 
readily in water. (Compare § 131.) The solution is hydro¬ 
fluoric acid. The gas and solution are dangerous substances. 
The gas is poisonous, and the acid if dropped on the skin 
produces terrible sores. Owing to its corrosive properties, 
hydrofluoric acid is kept in wax or bakelite bottles. 

Bakelite is a resinous substance, which resists the action 
of many chemicals, being practically inert. It has numerous 
applications, particularly in electrical apparatus. 

Hydrofluoric acid behaves chemically much like hydro¬ 
chloric acid. It is unlike hydrochloric acid in one respect; it 
forms both normal and acid salts, e.g., potassium fluoride 
(K 2 F 2 ) and acid potassium fluoride (KHF 2 ). 

431. Etching with hydrofluoric acid. — The acid and 
moist gas attack glass, and are used extensively in etching. 
Glass is essentially a mixture of silicates and silica (§§ 449, 
460) . Hydrofluoric acid interacts with these compounds and 
forms among other substances the volatile compound called 
silicon tetrafluoride (SiF 4 ). Hence the acid disintegrates the 
glass — literally “ eats ” or etches it. Typical equations for 
the reactions are : — 


334 


A BRIEF COURSE IN CHEMISTRY 


CaSiOs 3 H 2 F 2 = SiF4 CaF2 H“ 3 H 2 O 

Calcium Hydrofluoric Silicon Calcium Water 

Silicate Acid Tetrafluoride Fluoride 

Si02 “h 2 H 2 F 2 = SiF4 2 H 2 O 

Silicon Dioxide Hydrofluoric Acid Silicon Tetrafluoride Water 


In etching with hydrofluoric acid, the glass is thinly coated with wax, 
and the design or marks to be etched are scratched through the wax. 

The glass is then exposed to the gas or liquid, which 
attacks the unprotected places. When the wax is 
removed, a permanent etching is left. Hydrofluoric 
acid is used in marking the scales on thermometers, 
tubes, and other graduated glass instruments, and 
also in etching designs on glassware (Fig. 190). 

BROMINE 

432. Occurrence and preparation of bro¬ 
mine. — Bromine is never found free, but. 
bromides are widely 
distributed, especially 
magnesium and so¬ 
dium bromides. The 
salt springs and wells 
of Ohio, West Virginia, 

Pennsylvania, and 
Michigan contain bromides. Large 
quantities are found in the salt de¬ 
posits at Stassfurt in Germany. 

Bromine is prepared in the labo¬ 
ratory by heating a bromide with 
manganese dioxide and sulfuric acid 
(Fig. 191), thus: — 

2 KBr + 2 H2SO4 + Mn02 = Br2 + MnS 04 + K2SO4 + 2 H2O 

Potassium Sulfuric Manganese Bromine Manganese Potassium Water 

Bromide Acid Dioxide Sulfate Sulfate 



Fig. 191. —Apparatus for 
preparing bromine in the 
laboratory. 



Fig. 190. — Etch¬ 
ing on glass 
tumbler. (De¬ 
signed and exe¬ 
cuted by a high 
school pupil.) 


Another method consists in warming a bromide solution 
with chlorine; an equation for this method is : — 

MgBr2 + CI 2 = Br2 + MgCL 

Magnesium Bromide Chlorine Bromine Magnesium Chloride 










FLUORINE — BROMINE — IODINE 


335 


The chlorine method is used to manufacture bromine 
from the concentrated solution of bromides (largely magne¬ 
sium bromide) called bittern, which is left after sodium 
chloride has been removed by crystallization from the brine 
of salt wells. 


A sketch of the apparatus is shown in Figure 192. The bromide 
solution enters the tank at R, falls as a spray upon the resistant material, 
trickles down, and meets the ascending 
chlorine (gas) which is forced in at A. The 
liberated bromine escapes as a vapor 
through C, and the magnesium chloride 
flows out through D. In this process, chlo¬ 
rine, being the more active element, dis¬ 
places bromine, thus, 

CI 2 + 2 Br- —^ Bra + 2 Cl". 

433. Properties of bromine. — 

Bromine is a dark red liquid which 
is about three times as heavy as 
water. It is a volatile liquid, boil¬ 
ing at 59° C. The vapor, which has 
a disagreeable odor, irritates the 
mucous membrane of the eyes, nose, 
and throat. A bottle of bromine 
should not be opened unless it is in 
the hood. 

Bromine is moderately soluble in 
water. The solution is called bro¬ 
mine water. Bromine dissolves in 
carbon disulfide and carbon tetra¬ 
chloride ; the solution is reddish 
yellow. Liquid bromine burns the 
flesh frightfully, and care should be used in preparing it or 
working with it. 

The chemical behavior of bromine is similar to that of 
chlorine, though bromine is less active. It combines with 
metals, e.g., magnesium and iron. 

434. Compounds of bromine. — These are similar to those 
of chlorine. Hydrogen bromide (HBr) is a colorless, pun¬ 
gent gas, which fumes in the air and dissolves freely in water. 



Fig. 192. — Sketch of the 
apparatus for manufac¬ 
turing bromine. 











336 


A BRIEF COURSE IN CHEMISTRY 


This solution is called hydrobromic acid; it is much like 
hydrochloric acid. Bromides are salts of hydrobromic acid. 
Potassium bromide (KBr) is a white solid, made by decom¬ 
posing iron bromide with potassium carbonate; it is used as 
a medicine. Silver bromide (AgBr) is a pale yellow solid, 
and is used extensively in making photographic films and 
plates. ' 

IODINE 

435. Occurrence of iodine. — Iodine, like fluorine, chlorine, 
and bromine, is found in nature only in compounds. To¬ 
bacco, water cress, cod-liver oil, oysters, and marine plants 
contain minute quantities. Sea water contains a very small 
proportion of iodine compounds, which are assimilated by 
seaweeds. Iodine was formerly extracted from the ash of 
certain seaweeds. 

Iodine compounds, chiefly sodium iodate (NalOs), occur in 
the deposits of sodium nitrate in Chile, and most of the iodine 
of commerce is obtained from this source. 

436. Preparation of iodine. — Iodine is prepared in the 
laboratory by the same method as that used for bromine; 
potassium iodide is used instead of potassium bromide. 

Iodine is manufactured from sodium iodate which is left in the Chile 
sodium nitrate residues. The equation is: — 

2 NalOs + 3 NaaSOs + 2 NaHSOs =12 + 5 Na 2 S 04 + H 2 O 

Sodium Sodium Acid Sodium Iodine Sodium Water 

Iodate Sulfite Sulfite Sulfate 

437. Properties of iodine. — Iodine is a lustrous, gray- 
black crystalline solid. When gently heated, it changes 
into a beautiful violet-colored vapor, which quickly solidifies 
on a cold surface. This property of iodine, viz., ready 
transformation from solid into vapor and back directly 
into solid, is utilized in purifying iodine. The crude sub¬ 
stance is heated gently and the vapor condensed; the non¬ 
volatile impurities remain behind. This process is called 
sublimation and is frequently used to purify substances, e.g., 
ammonium chloride (§ 196). 

Iodine dissolves slightly in water but freely in alcohol, 
chloroform, carbon disulfide, ether, carbon tetrachloride, 


FLUORINE — BROMINE — IODINE 


337 


and potassium iodide solution. The chloroform and carbon 
disulfide solutions are violet, but the others are brown, or 
even black. Iodine and its solutions turn the skin brown. 

Iodine turns a cold starch suspension blue. The presence 
of starch in many vegetable substances can thus be readily 
shown (§ 261). A minute trace of iodine may be detected 
by starch, and this experiment also serves as a test for iodine. 
The colloidal starch adsorbs the iodine. 

In chemical properties iodine resembles chlorine and 
bromine, but iodine is less active. Bromine and chlorine 
displace iodine from many of its compounds. 

438. Uses of iodine. — A solution of iodine in alcohol 
(or in alcohol and potassium iodide), called tincture of iodine, 
or merely “ iodine,’^ is applied to the skin to harden it, to 
prevent the spread of eruptions, or to reduce swellings. 
Iodine is used to make iodoform (CHI 3 ), which is an anti¬ 
septic for wounds. Large quantities of iodine are made 
into iodides, drugs, and dyes. 

439. Compounds of iodine. — These are similar to the 
corresponding compounds of chlorine and bromine. Iodides 
are salts of hydriodic acid (HI). The best known salt is 
potassium iodide (KI). Silver iodide (Agl), like silver 
bromide, is used in photography (§ 491). 

440. The halogen elements and the periodic classification. 
— The halogen elements illustrate the periodic classification 
(§§ 418-421). These elements, as arranged in the periodic 
table, increase in atomic weight from fluorine through chlorine 
and bromine to iodine, and many of their properties are 
graded in this order. Thus, as we pass from fluorine to 
iodine the specific gravity increases, the color grows deeper, 
the volatility decreases, and the melting points of the solidi¬ 
fied elements increase. The intensity of the chemical action 
decreases as we pass from fluorine to iodine. 

EXERCISES 

1. Summarize the chief properties of fluorine and hydrogen fluoride. 

2. Describe the process of etching glass. Write the equations. 

3. Write the equations for the preparation of (a) hydrogen fluoride, 
(b) bromine, (c) iodine, (d) silver bromide. Write (d) in ionic form. 


338 


A BRIEF COURSE IN CHEMISTRY 


4. Write the formulas of the fluoride, bromide, and iodide of Al, 
ammonium, Ba, Ca, copper (ous and ic), Fe+^, Fe+^, Pb+^, magnesium, 
Sb+^ Si, Hg (ous and ic), Sn+2, Sn+^, zinc. 


PROBLEMS 

1. Calculate the per cent of bromine or iodine in (a) sodium bromide, 

(b) hydrogen bromide, (c) calcium iodide, (d) sodium iodate. 

2. How much (a) calcium sulfate and (b) hydrogen fluoride are 
formed by heating 60 gm. of fluor spar with sulfuric acid? 

3. How much potassium bromide (95 per cent pure) is necessary to 
prepare 47 gm. of bromine? 

4. Calculate the atomic weight of fluorine, bromine, or iodine from 
the following: (a) 1 gm. of CaF 2 gives 1.745 gm. of CaS 04 ; (6) 3.946 gm. 
of Ag (dissolved in HNO 3 ) require 4.353 gm. of KBr for precipitation; 

(c) 6.3835 gm. of silver iodide give 3.8965 gm. of silver chloride. 

6 . Calculate the simple formulas corresponding to (a) F = 48.72, 
Ca = 51.28; (b) Br = 67.22, K = 32.77; (c) I = 76.5, K = 23.49. 


SUGGESTIONS FOR LABORATORY WORK 


(References are to Newell’s Laboratory Exercises in Chemistry) 


Exercise 

Exercise 

Exercise 

Exercise 

Exercise 

Exercise 

Exercise 

Exercise 

Exercise 

Exercise 

Exercise 


544 — Preparation and Properties of Bromine — T. 

545 — Bromine (Short Method). 

546 — Tests for Free and Combined Bromine. 

S59 — Testing Salts (Bromide part only). 

S60 (e) — Silver Salts and Photography — T. 

547 — Preparation and Properties of Iodine — T. 

548 — Iodine (Short Method). 

549 — Tests for Free Iodine. 

550 — Tests for Combined Iodine. 

55 (c), (d) — Tests for Starch. 

S59 — Testing Salts (Iodide part only). 


TOPIC XXI: SILICON DIOXIDE — SILICATES — SILICON 
CARBIDE — SILICON TETRAFLUORIDE — GLASS 

441. Occurrence of silicon dioxide. — The most abundant 
and important compound of silicon is silicon dioxide (Si 02 ). 
It is often called silica. Sand, gravel, sandstone, and the 
numerous varieties of quartz are almost wholly silicon di¬ 
oxide; many rocks, e.g., granite and gneiss, contain silica 
as an essential ingredient. Ordinary soil contains more or 
less silica. 


SILICON 


339 


Quartz is the commonest form of silicon dioxide. Pure 
quartz is colorless and transparent, and is called rock crystal. 
There are many colored varieties, which are used as semi¬ 
precious jewels, e.g., amethyst, onyx, carnelian. It is crys¬ 
talline and is found in groups or as single crystals (Fig. 193). 



Fig. 193. — Quartz crystals—group (left), single crystals (center and right). 


Petrified or silicified wood is the remnant of wood in which 
the fiber has been replaced by silica. There is a forest of 
petrified trees in Arizona. Infusorial or diatomaceous earth 
(also called Tripoli powder) is largely silica and consists of 
the skeletons of minute organisms called diatoms, many 
being of delicate and beautiful structure (Fig. 194). 

442. Properties of silicon dioxide. — Most varieties of 
silicon dioxide are hard and brittle. Quartz is harder than 
other common sub¬ 
stances, and it breaks 
into fragments with 
sharp edges. 

Quartz melts at 

about 1600 C. and — Diatom shells (enlarged), 

can be melted only in 

the oxy-hydrogen fiame or the electric furnace. If pure silica 
is fused with certain precautions, the viscous mass can be 
drawn into elastic threads, which are used to suspend delicate 
parts of electric instruments; it can also be shaped into 
tubes, flasks, crucibles (Fig. 195), and even large pieces of 
apparatus used in industrial processes, e.g., dishes for crystal¬ 
lizing corrosive solutions and tubes for condensing acid 
vapors. 

By melting fragments of quartz (rock crystal) in an electric 














340 


A BRIEF COURSE IN CHEMISTRY 


furnace, first under reduced and finally under increased pres¬ 
sure, transparent quartz can be produced as rods, sheets, 
disks, lenses, and other forms, which are used in moving 
picture projection, photography, astronomy, and microscopy 
(Fig. 196). 

Objects made of fused quartz expand or contract only a 
very little within a wide range of temperature, i.e., they have 
a low coefficient of expansion, being only about 0.00000449 
between 0° C. and 1000° C. Thus, a silica crucible, unlike 
a porcelain one, will not crack if heated and then plunged at 
once into cold water. Similarly, the quartz lenses, which 
have largely replaced the glass condensing lenses in moving 



Fig. 195. — Apparatus made from fused silica for use in the laboratory. 


picture projection machines, do not crack even at the high 
temperature of the electric arc. 

Silicon dioxide does not dissolve in water. Nor is it at¬ 
tacked by acids, except hydrofluoric acid (§§ 430, 431). It 
is converted into a soluble silicate when boiled in water con¬ 
taining alkaline substances or when fused with the hydroxides 
or the carbonates of sodium and potassium, thus: — 

Si02 T Na2C03 = Na2Si03 -J- CO 2 

Silicon Dioxide Sodium Carbonate Sodium Silicate Carbon Dioxide 

443. Uses of silicon dioxide. — Sandstone is used as a 
building stone, and some varieties of sandstone are shaped 
into grindstones and whetstones. Large quantities of sand 
are used to make glass, porcelain, cement, and mortar. 
Different grades of sand are used as grinding and polishing 
material. Glass is polished by rubbing it with fine wet sand; 
it is also roughened and cut by blowing or blasting ” fine 








SILICON 341 

sand against it. Glass stoppers for bottles used in the labo¬ 
ratory are “ ground ” with sand. 

Infusorial earth (Fig. 194) is used as a polishing powder 
for metals electro-silicon ” being the commercial name 
of one kind), and in making scouring soaps. Owing to the 
hollow structure of its minute particles, considerable is used 
as an absorbent of nitro¬ 
glycerin in the manufac¬ 
ture of dynamite. 

444. Silicates. — 

These are salts of silicic 
acids. Only two silicic 
acids are well known, 
viz., metasilicic acid 
(H 2 Si 03 ) and orthosili- 
cic acid (H 4 Si 04 ). The 
salts of these and other 
sihcic acids (not yet 
prepared) are among the 
most common sub¬ 
stances in the earth’s 
crust. 

Most rocks consist 
wholly or largely of sili¬ 
cates of the metals 
aluminum, iron, calcium, 
potassium, sodium, and 
magnesium. Examples of silicates are slate, asbestos 
(Mg 3 Ca(Si 03 ) 4 ), feldspar (KAlSi308), mica, hornblende, 
clay (H 2 Al 2 (Si 04 )^. H 2 O), beryl, garnet, serpentine 
(Mg 3 Si 207.2 H 2 O), and talc (H 2 Mg 3 (Si 03 ) 4 ). The lava 
ejected by volcanoes consists largely of fused silicates. 

445. Sodium silicate. — When fine sand and sodium 
carbonate are fused together, sodium (meta-) silicate 
(Na 2 Si 03 ) is formed (see § 442 end). Sodium silicate is a 
glassy solid, which dissolves in water; hence this silicate is 
sometimes called water glass. Silicate of soda is the name 
given to commercial solutions of varying composition, which 
also contain some silica in the colloidal state (§§ 95, 446). 


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Fig. 196. — Photograph of a printed page 
made through disks of transparent quartz 
nine inches thick. 



342 


A BRIEF COURSE IN CHEMISTRY 


Solutions of sodium silicate have different properties 
adapted to special uses. Certain brands are used as the 
adhesive in making corrugated paper board, fiber board, 
and wall board. Paper board and fiber board are made 
into cartons, tubes, trunks, and containers of various kinds 
(Fig. 197). Wall board is used as a substitute for wood. 
Another brand of silicate of soda is used as a binder in 
abrasive wheels, stove and furnace cements, and linings. 
Still another use is as a protective coating for wood, paper. 



Fig. 197. — Making 4-ply paper board used for cartons, shipping con¬ 
tainers, and wall board. The two center strips are coated on both sides 
with sodium silicate solution; the upper and lower strips on one side 
only. All four strips are united firmly together by pressure rolls. This 
machine is one of the largest in use and turns out over 200 feet of finished 
board per minute. 

and concrete, a fireproof coating or ingredient of wood and 
fabrics, and a preservative coating for eggs. The last 
use is based on the fact that the deposit of silica in the 
pores of the shells keeps out the air and thus prevents 
the eggs from spoiling. Finally, considerable is used in 
many chemical processes, e.g., the manufacture of certain 
laundry soaps, sizing of paper, refining of oil, and removing 
paint. 

446. Colloidal silicic acid. — When hydrochloric acid is 
added to sodium silicate solution, a white gelatinous precipi¬ 
tate called silicic acid is formed, thus: — 






SILICON 


343 


Na 2 Si 03 + 2 HCl = HgSiOs + 2 NaCl 

Sodium Hydrochloric Silicic Sodium 

Silicate Acid Acid Chloride 

Under certain conditions probably the precipitate is a 
mixture of meta- and ortho-silicic acids (H 2 Si 03 and H 4 Si 04 ). 
These acids are closely related. The ortho-acid passes into 
the meta-acid upon drying, thus: — 

H4Si04 = H2Si03 + H 2 O 

Orthoailicic Acid Metasilicic Acid Water 

And metasilicic acid when heated decomposes into silicon 
dioxide and water, thus: — 

H2Si03 = Si02 + H 2 O 

Metasilicic Acid Silicon Dioxide Water 

Silicon dioxide may be regarded as the anhydride of meta¬ 
silicic acid, though contrary to the usual rule this acid can 
not be formed from the anhydride (Si 02 ) and water (§§86, 
219). 

Sodium silicate and hydrochloric acid do not always inter¬ 
act as just described. That is, the silicic acid is not gelati¬ 
nous, but is dispersed throughout the liquid in the colloidal 
state (§ 95). This colloidal silicic acid can not be filtered 
out because the colloidal particles are so fine they pass 
through the paper. It can be separated by dialysis from the 
sodium chloride in the solution. Thus, if the colloidal sus¬ 
pension of silicic acid is placed in a vessel, which has a bottom 
of parchment, and hung in a larger receptacle filled with 
water, the sodium chloride will pass through the parchment 
into the water, but the silicic acid will be retained in the 
smaller vessel. 

The gelatinous silicic acid is an example of a hydrogel, or, 
more frequently called, a gel. Whereas the coUoidally sus¬ 
pended silicic acid is an example of a hydrosol, or simply a 
sol. 

447. Silicon carbide or carborundum. — This is a crys¬ 
tallized solid, which varies in color from white to green or 
black. It is extremely hard, being nearly as hard as diamond. 
It is very resistant to chemicals. The extreme hardness of 


344 


A BRIEF COURSE IN CHEMISTRY 



carborundum has led to its extensive application as an 
abrasive, and large quantities are made into a great variety 
of grinding wheels, whetstones, and polishing cloths (Fig. 18). 

Carborundum is manufactured by fusing a mixture of sand (silicon 
dioxide, Si02) and coke (carbon, C) in an electric furnace similar in form 
to that used in the manufacture of graphite (Fig. 198). The mixture of 
sand and coke (to which salt and sawdust are added to contribute to 
the fusion and porosity) is packed around the core of granulated coke. 
The heat generated by the resistance of the carbon core to the passage 


Fig. 198. — Furnace for the manufacture of carborundum. 

of the powerful current of electricity produces a chemical change essen¬ 
tially as follows: — 

Si02 + 3 C = SiC + 2 CO 

Silicon Dioxide Carbon Silicon Carbide Carbon Monoxide 

448. Silicon tetrafluoride. — When hydrofluoric acid inter¬ 
acts with silicon dioxide, silicates, or silicon, silicon tetra¬ 
fluoride (SiF 4 ) is formed (§ 431). Silicon tetrafluoride is a 
colorless, volatile liquid. In the World War it was used to 
produce a smoke screen, because with ammonia and water 
it forms a dense white smoke consisting of silicic acid and 
ammonium chloride (Fig. 56). It fumes in moist air owing to 
interaction with water. The equation for the reaction is: — 

3 SiF4 + 4 H 2 O = H4Si04 + 2 HsSiFe 

Silicon Tetrafluoride Water Silicic Acid Hydrofluosilicic Acid 




SILICON 


345 


The hydrofluosilicic acid (sometimes called fluosilicic acid) 
remains in solution, while the (ortho) silicic acid is precipi¬ 
tated. The formation of the white gelatinous silicic acid 
when the gases from the interaction of hydrofluoric acid and 
a compound of silicon are led into water is often used as a 
test for silicon. 

449. Glass. — This substance is essentially a mixture of 
silicates and silica. Ordinary glass, such as that used for 
bottles, is made by heating sand, calcium carbonate, and 
sodium carbonate to a high temperature. The product is 
a mixture of calcium silicate (CaSiOs), sodium silicate 
(Na 2 Si 03 ), and silica (Si02). Sodium-calcium glass is called 
soft glass because it softens readily when heated. 

Other varieties can be made by substituting other sub¬ 
stances, wholly or partly, for the calcium and sodium com¬ 
pounds. Thus, potassium carbonate (with calcium car¬ 
bonate) produces a hard glass, which melts at a higher 
temperature than sodium glass. 

If a lead compound is used in place of calcium carbonate, 
flint glass is produced which refracts light to a high degree; 
it is used to make lenses for optical instruments, and shades 
for gas and electric lights. On account of its brilliancy this 
glass i^ made into cut glass vessels for ornaments and table 
use. 

Another kind of glass, used extensively for chemical ap¬ 
paratus and certain cooking utensils, contains a large excess 
of silica together with boron oxide (B2O3) and aluminum 
oxide (AI2O3). It is made by heating together sand, borax 
(Na 2 B 407 ), and aluminum oxide. This glass is tough and 
has a much lower coefficient of expansion than ordinary 
glass; hence it does not crack with sudden changes of tem¬ 
perature. It is called pyrex glass. 

460. Manufacture of glass. — The ingredients needed for the 
different varieties of glass are mixed in the proper proportions and heated 
to a high temperature in a long tank, or in large clay vessels. The sub¬ 
stances interact and form a heavy, viscous, semi-fluid mass, which can be 
poured into molds, rolled into sheets, and fashioned into various shapes 
by blowing, pulling, pressing, or stamping. 

Bottles, jars, and electric light bulbs are manufactured by complicated 
machines. Special or ornamental objects are blown into a mold. A 


346 


A BRIEF COURSE IN CHEMISTRY 



Fig. 199. — Blowing huge cylinders of glass by compressed air. These 
are slit open and allowed to roll out flat. 


skilled workman gathers a mass of the plastic glass on the end of a long 
tube, called a glass blowpipe, blows the glass into a preliminary shape, 
lowers it into a mold, and blows until the mold is filled. 

Window glass and other kinds of flat glass are made by machinery. 
(1) In one process the glass is blown by compressed air into a long 
cylinder (Fig. 199) which is allowed to cool; the ends are then cut off, 
the cylinder is slit lengthwise, heated until it softens and opens, and 
then allowed to roll out flat. (2) In another process the mixture of 
the raw materials is put into a furnace, melted, and transformed into 








SILICON 


347 



Fig. 200. — Removing a sheet of glass from an automatic sheet glass ma¬ 
chine. The raw materials enter at one end and the sheets of glass are 
cut off at the other. 

glass. The viscous mass is moved along (in the same machine) to a bar 
which sticks to it, pulls it forward between rolls, through a cooling cham¬ 
ber, and finally out upon a table where the flat glass is automatically 
cut into sheets. In Fig. 200 a workman is shown taking off a sheet of 
glass. 

Plate glass, which is used for windows and wind shields, is made by 
the flat process or by pouring the molten glass into a shallow frame upon 
a large table, rolling it with a hot steel roller, as if it were dough, and 



Fig. 201. — Bottles just coming from an annealing chamber which is called 

a lehr. 






















348 


A BRIEF COURSE IN CHEMISTRY 


finally grinding and polishing the cooled sheet on a rotating table until 
the surfaces are parallel. 

Glass must be cooled slowly, otherwise it will crack or crumble to 
pieces when jarred or scratched. This slow cooling is called annealing. 
It is accomphshed by passing the objects on a slowly moving frame 
through a long chamber in which the temperature is gradually lowered 
(Fig. 201). 


EXERCISES 

1. Summarize the properties of silica. 

2. Describe the formation, state the uses^ and enumerate the prop¬ 
erties of water glass. 


PROBLEMS 

1. Calculate the per cent of silicon in (a) orthosihcic acid, (6) meta- 
silicic acid, (c) potassium feldspar (KAlSiaOg). 

2. How much sodium silicate can be made from 1000 kg. of sand? 

3. Calculate the simplest formulas corresponding to (a) Si = 35.897, 
H = 2.564, O = 61.538; (b) Si = 29.166, H = 4.166, O = 66.666. 


TOPIC XXH: COPPER 

451. History. — Copper has been known for ages. Do¬ 
mestic utensils and weapons of war containing copper were 
used before similar objects of iron. The Latin word cuprum 
gives the symbol Cu. 

452. Copper ores. — Free or native copper, mixed with 
a hard rock, is found in large quantities in the Michigan 
mines on the shore of Lake Superior. The most abundant 
copper ores are the copper-iron sulfide ores, e.g., chalcopyrite 
(CuFeS 2 ), and bornite (CusFeSs), found in Montana and 
LFtah (Fig. 202). 

453. Metallurgy of copper. — Free copper is easily ex¬ 
tracted. The ore is crushed, concentrated by grinding and 
then washing away the rocky impurities down an inclined 
plane or on shaking tables, and finally heated (Le., smelted) 
with a flux. 

The metallurgy of the copper-iron sulfide ores is compli¬ 
cated. It involves the removal of sulfur and iron. The 
result is accomplished by converting the sulfur into sulfur 


COPPER 


349 



dioxide, which escapes as 
a gas (or is made into 
sulfuric acid), and the 
iron into ferrous silicate, 
which is removed as slag. 


464. Metallurgy of cop¬ 
per-iron-sulfide ores. — Let us 

consider six steps in detail by 
using Fig. 203. 

1. Crushing. The ore is 
first crushed to the proper 
size in A. 

2. Concentrating. The 
crushed ore is next concen¬ 
trated by grinding, washing, 
shaking, settling, and floating. 

By these operations useless 
parts of the ore are removed 
and three general sizes result: 

(a) The coarse concentrate, 
which goes from the shaking 
machine B to the blast fur¬ 
nace M. (See 4 (2).) (6) The 
fine concentrate, which, after 
the grinding in C and shaking 
in D, goes to the roasting fur¬ 
nace K. (c) The very fine concentrate, which is further concentrated. 
The ground ore from F passes through the desliming cone F on to the 


Fig. 202. — Workman in a Montana 
copper mine. 



Fig. 203. — Diagram showing the steps in obtaining copper from copper- 

iron-sulfide ores. 












































































350 


A BRIEF COURSE IN CHEMISTRY 


table G; here part goes to the roasting furnace K. The rest is ground 
exceedingly fine in the mill H and goes on into the flotation machine I-J 
(Fig. 204). Here a remarkable change occurs. The fine ore particles 

are mixed with water contain¬ 
ing a little oil (and sometimes 
sulfuric acid). Air is beaten 
into the mixture by vigorous 
agitation. An oily froth is 
formed to which the particles 
of copper ore stick and float on 
the top of the mass, whereas 
the rocky particles sink. The 
water is removed (by a filter 
— not shown) and the con¬ 
centrate goes to the roasting 
furnace K. 

3. Roasting. The charge is 
heated red-hot but not melted; 
the heat, after the furnace is 
once started, comes from the 
burning sulfur. By this treat¬ 
ment most of the sulfur is 
removed and part of the copper and iron sulfides are changed into 
oxides. The product goes to the reverberatory furnace L. 

4. Smelting. The product from the roasting furnace K is smelted 
in the reverberatory furnace L; and the coarse concentrate from the 
jig R is smelted in the blast fur¬ 
nace M. That is, they are 
heated with a flux until they 
melt, just as iron is smelted. 

By this treatment more sulfur 
(as sulfur dioxide) and iron (as 
slag) are removed, and the 
copper and iron sulfides melt 
together to form copper matte. 

There are two kinds of furnaces 
for smelting: 

(1) In the reverberatory fur¬ 
nace L the heat radiated down 
upon the hearth fuses the charge 
(Fig. 205). Most of the iron 
sulfide becomes iron oxide and 
forms a slag with the lime, silica, 
and alumina (AI2O3) in the 
charge, while the copper sulfide and the remaining iron sulfide melt 
together and sink through the slag to the bottom of the mass. The 
slag runs off continually. The matte is tapped off periodically and taken 
to the converter N. 



Fig. 205. — Reverberatory furnace. The 
fire burns on the grate G, and the long 
flame which passes over the bridge, 
E, is reflected down by the sloping 
roof upon the contents of the furnace. 
Gases escape through I. The charge, 
which rests upon B, does not come in 
contact with the fuel. 















































COPPER 


351 



(2) The blast furnace M is much like that used in making cast iron 
(Figs. 126, 127, 128), though it is cooled by a constant flow of water 
through jackets which inclose its sides. The charge consists of the 
coarse concentrate (see 2 (a) above), limestone, and coke. Air is blasted 
through the furnace. Much heat is supplied by the burning sulfur. 
By this treatment much of the sulfur is removed and most of the iron 
forms a slag with the hmestone in the charge or with the silica and 
alumina purposely left in the concentrate. The copper sulfide melts 
with the rest of the iron sulfide into copper matte, which goes to the 
converter N. 

5. Converting. This is the last stage. Matte from the reverberatory 
and blast furnaces (L and M) is poured into the converter — 65 tons to 


Fig. 206. — A copper converter in action. 

a charge (Fig. 206). The converter is lined with magnesia brick. Air 
is blown through the liquid mass. The remaining sulfur burns to sulfur 
dioxide which escapes through the top, while the iron forms a slag with 
the sihca and alumina in the ore (or added to supply these substances). 
The product, which is metallic copper together with a little iron, sulfur, 
and slag, goes to the refining-casting furnace O. 

6. Casting into anodes. The product from the converter N is further 
purified in the refining-casting furnace 0 by blowing air through it and 
then stirring it with poles of green wood to reduce any oxide. When 
the refining is complete, the copper, now 99.25 per cent pure and called 
blister copper, is cast into anodes. The anodes weigh about 500 pounds 
each. They are sent to the refinery for final purification by electrolysis. 

466. Refining of copper by electrolysis. — Since very pure 
copper (at least 99.95 per cent) is needed in electrical indus- 





352 


A BRIEF COURSE IN CHEMISTRY 


tries, the blister copper must be further purified to remove the 
last traces of silver, gold, and other impurities. This is 
done by electrolysis, and the refined metal, which is 99.98 per 
cent pure, is called electrolytic copper. 

The anodes are suspended in a solution of copper sulfate 
and sulfuric acid. The cathodes, which are made of thin 
sheets of pure copper (coated with graphite), also dip into the 
solution between the anodes (Fig. 207). When the current 
passes, copper ions (Cu++) migrate to the cathode, lose their 
charges, and are deposited as metallic copper, thus building 
up the cathode (Fig. 208). An equal weight of copper dis- 

solves from the anode, which 


n 


n 


n 

X 




Fig. 207. — Sketch of the apparatus 
for the preparation of pure copper by 
electrolysis. A, A, A are anodes, 
and C, C, C are cathodes. 


gradually wears away. Only 
copper is deposited on the 
cathode. The gold and 
silver and other substances 
drop to the bottom of the 
cell ^as a slime; from this 
slime the gold, silver, and 
other substances which were 
left in the copper anodes are 
profitably extracted. 

456. Properties of copper. — Copper is distinguishable 
from all other metals by its peculiar reddish color. It is 
flexible, ductile, malleable, and tough, and hence can be 
shaped into many forms. It melts at 1083° C. Its specific 
gravity is 8.9. Copper is an excellent conductor of elec¬ 
tricity — the best of the cheaper metals. 

Exposed to ordinary air, it turns dull owing to a thin 
film of oxide. In moist air it gradually becomes coated with 
a green (basic) copper carbonate. Heated in the air, it is 
first changed into black copper oxide (CuO), and at a high 
temperature it burns, and colors a flame emerald-green. 

Copper, unlike other common metals, e.g., magnesium, 
zinc, and iron, does not liberate hydrogen from acids (§§ 68, 
72). With nitric acid it forms copper nitrate and nitrogen 
oxides (§ 203) ; with hot sulfuric acid it yields copper sul¬ 
fate and sulfur dioxide (§ 217). Hydrochloric acid has 
little effect upon it. Sea water attacks it, forming a (basic) 







































COPPER 353 

cupric chloride; hence the copper sheathing of vessels cor¬ 
rodes. 

Copper displaces certain metals from their solutions 
(§ 462). A clean copper wire, if placed in a solution of any 
mercury compound, soon becomes coated with mercury. 
On the other hand, metals like iron, zinc, and magnesium 
displace copper from its solution, e.g., a nail or a knife blade 
soon becomes coated with copper if dipped into a solution of 
any copper compound. Scrap iron is used to precipitate 
copper on a large scale in some copper refineries. 



Fig. 208. — Removing cathodes of pure copper from a cell in one of the 
rooms of an electrolytic refining plant. 


467. Tests for copper. — (1) The reddish color, peculiar 
“ coppery ’’ taste, and green color imparted to a fiame 
serve to identify metallic copper. (2) An excess of ammo¬ 
nium hydroxide added to a solution of a copper compound 
produces a beautiful deep blue solution. (3) A few drops of 
acetic acid and potassium ferrocyanide (K 4 Fe(CN) 6 ) added 
to a dilute solution of a copper compound precipitate brown 
copper ferrocyanide (Cu 2 Fe(CN) 6 ). 

468. Uses of copper. — Large quantities of copper wire 
are used to conduct electricity, e.g., in operating the tele¬ 
graph, cable, telephone, radio, electric railway, and electric 
light. Sheet copper is made into household utensils, boilers, 






354 


A BRIEF COURSE IN CHEMISTRY 


and stills, and is also used for roofs and spouts. Most nations 
use copper as the chief ingredient of small coins. Much 
copper is utilized in electrical and other apparatus, especially 
now that copper can be cast without blowholes. Books are 
printed and illustrated from electrotype plates made by 
depositing copper upon an impression of the type or design 
in wax. 

Copper alloys. — Copper is an ingredient of many common 
and useful alloys, e.g., bronze, brass, and monel metal 
(Fig. 209). (See Table XXL) 

TABLE XXI. — Copper Alloys 


Name 

Copper 

Zinc 

Alu¬ 

minum 

Silver 
AND Gold 

Nickel 

Tin 

Aluminum bronze 

90-98 

_ 

2-10 

_ 

_ 

__ 

Bell metal .... 

78 

— 

— 

— 

— 

22 

Brass. 

63-93 

5-40 

— 

— 

— 

— 

Bronze. 

60-95 

25 

— 

— 

— 

3-8 

German silver . . . 

50-60 

— 

— 

— 

6-20 

— 

Gun metal .... 

90 

— 

— 

— 

— 

10 

Gold coin .... 

10 

— 

— 

Gold 90 

— 

— 

Monel metal . . . 

28 

— 

— 

— 

69 

— 

Nickel coin .... 

75 

— 

— 

— 

25 

— 

Silver coin .... 

10 

• — 

— 

Silver 90 

— 

— 


469. Cuprous and cupric compounds. — Copper forms 
two series of compounds — the cuprous and the cupric. 
The valence of copper is + 1 in cuprous compounds and + 2 
in cupric. 

Cupric salts are more common. All dilute solutions of 
ordinary cupric salts are blue owing to the presence of the 
blue cupric ion (Cu++). 

Soluble copper compounds are more or less poisonous. 
Cooking utensils made of copper should be used with care. 
Vegetables, acid fruits, and preserves, if boiled in them, should 
be removed as soon as cooked. The vessels themselves 
should be kept bright to prevent the formation of copper salts, 
which might contaminate the contents. Certain lower forms 

















COPPER 


355 



Fig. 209. — Monel metal traffic markers (white spots) on a crowded “loop” 
in Chicago. 


of plant life (algae) are poisoned by copper salts. Copper 
sulfate is sometimes added to ponds and reservoirs to destroy 
such growths, and thereby render the water suitable for 
drinking. 

460. Copper sulfate or cupric sulfate. — This is a blue 
solid. It is also called blue vitriol or bluestone. Copper 
sulfate solutions have an acid reaction owing to hydrolysis 
(§ 281). (See also § 86.) 

Copper sulfate is used in electric batteries {e.g., the gravity 
cell), in making other copper salts, and in copper plating and 
electrotyping. A mixture of copper sulfate and milk of lime, 
called Bordeaux mixture, is sprayed upon trees to destroy 
fungi and kill insects. 

Copper sulfate is prepared by treating copper scrap in the 
air with warm, dilute sulfuric acid or by oxidizing copper sul¬ 
fide. Some of the copper sulfate of commerce is a by-product 
in refining gold and silver with sulfuric acid. 





356 


A BRIEF COURSE IN CHEMISTRY 


461. Other copper compounds. — Cuprous oxide (CU 2 O) 
is precipitated as a reddish powder by heating Fehling’s 
solution a mixture of solutions of copper sulfate, sodium 
citrate, and sodium carbonate) with glucose; its formation 
serves as a test for glucose and sugars like it (§ 258). Cupric 
oxide (CuO) is a black solid formed by heating copper in air. 
It is used to remove sulfur compounds from petroleum. Cop¬ 
per nitrate (Cu(N 03 ) 2 ) is a blue, crystallized solid, formed 
by the interaction of copper and dilute nitric acid. It is a 
cupric salt. Copper bromide (CuBr 2 ) is a black crystallized 
solid, which in dilute solutions is blue (due to Cu++) and in 
concentrated solution is brown (due to CuBr 2 ). 

462. Displacement of metals. — We have already seen 
that iron, zinc, and other metals displace copper from copper 
salt solutions, and that copper itself displaces mercury 
(§ 466). The deposition of metallic copper and of mercury 
are examples of a chemical change in which most metals 
can participate. It has been found that the metals can be 
arranged in the order in which one metal can displace another 
from its solution. This order is sometimes called the dis¬ 
placement series. The arrangement of the common metals 
is shown in the accompanying list (compare § 72). In 
this series each free metal displaces succeeding metals from 
their solutions. (Read horizontally.) 

Displacement Series of the Common Metals 

Magnesium Aluminum Zinc Iron Tin Lead 

Hydrogen Copper Mercury Silver Platinum Gold 

An ionic equation for an example of displacement is: — 
Zn + Cu++ + SO4 = Cu + Zn++ + SO4 — 

Hydrogen is not a metal in the common acceptance of this 
term. But it is usually included in the displacement series 
of metals, because the metals that precede hydrogen displace 
it from most acids (§ 68). That is, hydrogen behaves in this 
respect like the members of the series. Recall that hydrogen 
is like metals in that it forms positive ions (H+). 


COPPER 


357 


EXERCISES 

1. Prepare a summary of the metallurgy of copper-iron sulfide ores. 

2. Interpret the electrolytic refining of copper. 

3. State physical properties of copper that fit it for electrical uses. 

4. Give three tests for copper. 

6. Write the formulas of the following compounds by applying the 
principle of valence: Cupric bromide, cuprous chloride, cupric phos¬ 
phate, cupric sulfate, cuprous iodide. 


PROBLEMS 


1 . Calculate the weight of copper in 10 gm. of crystallized copper 
sulfate. 

2. Calculate the simplest formulas corresponding to: (a) Cu = 
96.94, H = 3.05; (b) Cu = 57.46, H = 0.91, O = 36.2, C = 5.43. 

3. How much copper can be obtained from an American ten cent 
coin which weighs 2.44 gm. ? 

4. A certain weight of copper oxide, when heated in a current of 
hydrogen, lost 59.789 gm. of oxygen and formed 237.55 gm. of copper. 
If the atomic weight of oxygen is 16, calculate the atomic weight of 
copper. 

SUGGESTIONS FOR LABORATORY WORK 


(References are to Newell’s Laboratory Exercises in Chemistry) 


Exercise S4 —Preparation of Copper Oxide. 

Exercise *6 — Reduction of Copper Oxide by Carbon. 

Exercise *10 — Reduction of Copper Oxide by Illuminating Gas. 

Exercise SIO — Reduction of Copper Oxide by Hydrogen — T. 

Exercise 12 — Test for Copper ((1) part only). 

Exercise *14 I — Solubility of Solids. 

Exercise *18 A — Water of Hydration (Copper Sulfate). 

Exercise S13 — Anhydrous Compounds (Copper Sulfate). 

Exercise SI5 — Efflorescence and Deliquescence (Copper Sulfate 

part only). 

Exercise 22 — Formation of Copper Sulfide. 

Exercise 33 — Electrolysis of Copper Sulfate Solution — T. 

Exercise S31 — Electrolysis of Copper Sulfate Solution (Short 

Method) — T. 

Exercise 35a — Colored and Colorless Ions. 

Exercise 34 — Hydrolysis of Certain Salts (Copper Sulfate only). 

Exercise *39 — Nitric Acid and Copper (III only). 

Exercise S38 — Sulfides (Copper part of (a)). 

Exercise 54 — Cuprous Oxide. 

Exercise 62c — Flame Tests for Metals. 

Exercise S56 B — Tests for Metals (Copper). 


358 


A BRIEF COURSE IN CHEMISTRY 


TOPIC XXIII: MAGNESIUM —ZINC —MERCURY 

463. Physical and chemical properties of magnesium. — 

Magnesium is a lustrous, silvery white metal. It is light, 
the specific gravity being 1.75. It is tenacious and ductile, 
and when hot can be drawn into wire or rolled into a ribbon, 
the latter being a common commercial form. 

Heated in air, it burns with a dazzling light and produces 
magnesium oxide (MgO) together with a little magnesium 
nitride (Mg 3 N 2 ). It does not tarnish in dry air, but in 



Fig. 210. — Pipes covered with “85 per cent magnesia” to prevent loss 
of heat. 


moist air it is soon covered with a film of the (basic) car¬ 
bonate. It liberates hydrogen from acids (§ 72). It also 
liberates hydrogen, slowly, from boiling water (§ 69). Solu¬ 
tions of magnesium salts contain magnesium ions (Mg++). 

464. Uses of magnesium. — The light from burning mag¬ 
nesium affects a photographic plate, and magnesium powder 
(mixed with potassium chlorate) is used in taking flashlight 
photographs. The powder is one ingredient of the mixture 
used in signal lights (e.gr., star shells) and fireworks. Magne¬ 
sium and its alloys are used to construct articles needing 
a light durable framework. Magnalium is an alloy of mag¬ 
nesium and aluminum (§ 604). 















MAGNESIUM — ZINC — MERCURY 


359 


465. Magnesium oxide and hydroxide. — Magnesium 
oxide (MgO) is a white, bulky powder. It is formed when 
magnesium burns in the air, e.g., in taking a photograph with 
a flashlight powder. It is manufactured by gently heating 
magnesium carbonate, somewhat as lime is made from lime¬ 
stone (§ 293). It is called magnesia, or calcined magnesia, 
though commercial magnesia is a complex compound (largely 
Mg(OH) 2 .3 MgCOs). Like lime, magnesia is infusible and 
is used in making Are brick, crucibles, and furnace linings. 
It is a poor conductor of heat, and is the main ingredient of a 
mixture called “85 per cent magnesia,’’ which is used to 
insulate pipes and boilers 



and thereby prevent the 
loss of heat (Figs. 210, 
211 ). 


Magnesium oxide 
unites slowly with water 
and forms magnesium 


Fig. 211. —Section of a pipe showing 
end of protective cover. 


hydroxide (Mg (OH) 2 ), 

which is a white solid. It is only very slightly soluble in 
water, and a suspension, called milk of magnesia, is used as a 
medicine to neutralize acidity in the stomach. Magnesium 
hydroxide is also an ingredient of some kinds of tooth paste. 

466. Magnesium sulfate and chloride. — Magnesium sul¬ 
fate (MgS 04 ) is a white solid. The commercial form is often 
called Epsom salts (MgS 04 .7 H 2 O). It is very soluble in 
water and its solution has a bitter taste; it is used as 
a purgative. Magnesium sulfate, like calcium sulfate, 
causes permanent hardness in water (§ 298). Magnesium 
chloride (MgCL) is a white solid. The crystallized salt 
(MgCL. 6 H 2 O) is very deliquescent. It is one of the 
ingredients which make common salt moist. Magnesium 
chloride causes permanent hardness in water. If water con¬ 
taining magnesium chloride, e.g., sea water, is used in a boiler, 
magnesium hydroxide and hydrochloric acid are formed. 
The insoluble magnesium hydroxide becomes a hard scale 
and the hydrochloric acid corrodes the metal. Hence water 
containing magnesium chloride (or sulfate) is highly objec¬ 
tionable and should be softened before use. 




360 


A BRIEF COURSE IN CHEMISTRY 


ZINC 

467. Occurrence of zinc. — The chief ores are zinc sulfide 
(sphalerite, zinc blende, ZnS) and zinc carbonate (smith- 
sonite, ZnCOs). 

468. Metallurgy of zinc. — The ores are usually concen¬ 
trated (sulfide ore by the flotation process (§ 453, 2 (c))), 
then roasted to form the oxide, which is reduced by heating 
with finely powdered coal. 

The reduction is conducted in special shaped earthenware retorts 
(A) connected with double receivers (Fig. 212), At first the zinc con¬ 
denses in C as a powder called zinc dust. But when this receiver be¬ 
comes hot, the zinc condenses to a liquid in B, from which it is drawn 
off at intervals and cast into bars or plates. The impure zinc, called 


Fig, 212. — Retorts for reduction of zinc oxide (open — right, closed — left). 



spelter, is freed from carbon, lead, iron, cadmium, and arsenic by re¬ 
peated distillation; very pure zinc is obtained by the electrolysis of a 
zinc salt (Fig. 213). 

469. Physical and chemical properties of zinc. — Zinc is a 
bluish white, lustrous metal. At ordinary temperatures it is 
rather brittle, but at 100^150° C. it is soft and can be rolled 
into sheets and drawn into wire; zinc which has been rolled 
or drawn does not become brittle on cooling. It melts at 
419° C. If melted zinc is poured into water, it forms irregu¬ 
lar brittle lumps called granulated zinc, which is a convenient 
form for use in the laboratory. 

Heated in air above its melting point, zinc burns with a 
bluish green flame, forming white zinc oxide (ZnO). Zinc 
does not tarnish in dry air, but ordinarily it becomes coated 
with a thin, non-porous film of basic zinc carbonate. Zinc 
interacts readily with acids and usually liberates hydrogen 
(§ 68, but see §§ 200, 203). With hot solutions of sodium 






MAGNESIUM — ZINC — MERCURY 


361 


and potassium hydroxides, it forms zincates and hydrogen, 
thus: — 

Zn + 2 KOH = K2Zn02 + H 2 

Zinc Potassium Hydroxide Potassium Zincate Hydrogen 

Zinc displaces most metals from their solutions (§ 462). 
Ordinary zinc salts yield zinc ions (Zn++) in solutions. 

470. Uses of zinc. — Zinc is extensively used as an elec¬ 
trode in many kinds of batteries. Sheet zinc is used for 
roofs, gutters, pipes, 
parts of washing ma¬ 
chines, and as a lining 
for tanks. 

The chief use of zinc is 
in making galvanized 
iron. This is iron cov¬ 
ered with a thin layer of 
zinc and is made by dip¬ 
ping clean sheet iron into 
melted zinc. The zinc 
protects the iron from 
air and moisture. Gal¬ 
vanized iron, if properly 
manufactured, does not 
rust easily and is exten¬ 
sively used for netting, 
wire, roofs, pipes, cor¬ 
nices, and water tanks. 

Zinc shavings are used 
in the cyanide process of 
extracting gold (§§ 494, 

496). Zinc is an ingredient of many useful alloys, e.g., brass, 
bronze, and German silver (§ 468). 

471. Zinc oxide. — The pure oxide (ZnO) is white when 
cold and yellow when hot. It is formed when zinc burns, 
and is manufactured in this way or by heating zinc carbonate. 
It is often called zinc white or Chinese white, and large quan¬ 
tities are used as a filler in the manufacture of automobile 
tires and white rubber goods. 



Fig. 213. — Lifting the cathodes of pure 
zinc from an electrolytic cell. 











362 


A BRIEF COURSE IN CHEMISTRY 


Zinc oxide is also used to make white paint. Paint made 
of zinc oxide, unlike paint made of lead compounds, is not 
darkened by sulfur compounds in the atmosphere (§ 620). 

On account of its antiseptic and drying properties, zinc 
oxide is an ingredient of ointments. 

472. Zinc hydroxide. — This is formed as a dull white 
precipitate by the interaction of a small amount of sodium 
or potassium hydroxide and a solution of a zinc salt. An 
excess of the alkaline hydroxide changes zinc hydroxide into 
a zincate (§ 469, end). Zinc hydroxide has both acid and 
basic properties. It dissolves in ammonium hydroxide 
owing to the formation of a soluble complex compound called 
zinc ammonia (or ammonio-) hydroxide (Zn(NH 3 ) 4 (OH) 2 ). 

473. Zinc sulfide, sulfate, and chloride. — Pure zinc sulfide 
(ZnS) is white, and is formed as a jelly-like precipitate when 
hydrogen sulfide is passed into a solution of a zinc salt. A 
mixture of zinc sulfide and barium sulfate, called Hthopone, is 
used as a white pigment in paints for interior work. 

The sulfate (ZnS 04 ) is formed by the interaction of zinc 
and dilute sulfuric acid. Large quantities are also made by 
roasting the sulfide in a limited supply of oxygen and extract¬ 
ing the sulfate with water. Thus prepared, it is a white, 
crystallized solid (ZnS 04 . 7 H 2 O) called white vitriol. It is 
used in dyeing, as a disinfectant, and as a medicine. Like 
other zinc salts, it is poisonous. 

The chloride (ZnCL) is a white, deliquescent solid. Large 
quantities are used to preserve wood, especially posts and 
railroad ties. 

474. Tests for zinc. — The precipitation of the sulfide or 
hydroxide, as above described, serves as a test for zinc. A 
green incrustation is produced when zinc compounds are 
heated on charcoal and then moistened with cobaltous nitrate 
solution. (Compare § 606, last paragraph.) 

MERCURY 

476. Metallurgy of mercury. — Mercury is prepared by 
roasting cinnabar (HgS) in a current of air, and condensing 
the vapor of the metal. 


MAGNESIUM — ZINC — MERCURY 


363 


476. Physical and chemical properties of mercury. — 

Mercury is a silvery metal, and is the only common one that 
is liquid at ordinary temperatures. Its common name 

quicksilver ” emphasizes these properties. It solidifies at 
— 38.7° C. and boils at 357° C. It is a heavy metal, the 
specific gravity being about 13.6. Mercury is a good con¬ 
ductor of electricity. 

Mercury does not tarnish in the air, unless sulfur com¬ 
pounds are present. At about 300° C. it combines slowly 
with oxygen to form red mercuric oxide (HgO). Hydro¬ 
chloric acid and cold sulfuric acid do not affect it; hot con¬ 
centrated sulfuric acid changes it into mercuric sulfate 
(HgS 04 ). Nitric acid changes it into nitrates — hot acid 
into mercuric nitrate (Hg(N 03 ) 2 ) and cold dilute acid, with 
an excess of mercury, into mercurous nitrate (HgNOa). It 
is displaced from solution by many metals (§ 462), and also 
displaces some, e.g., a copper wire becomes coated with 
mercury when put into a solution of a mercury compound. 

477. Amalgams. — Alloys of which mercury is a compo¬ 
nent are called amalgams. Amalgamated zinc is used in 
certain electric batteries to prevent unnecessary loss of the 
zinc. Amalgams of some metals {e.g., tin, silver, gold) are 
used as a filling for teeth. Silver and gold form amalgams 
readily, and considerable mercury is used in extracting these 
precious metals from their ores (§§ 483, 494). Care should 
be taken, while using mercury, not to let it come in contact 
with jewelry. 

478. Uses of mercury. — Mercury is used in thermom¬ 
eters, barometers, mercury-vapor lamps and illuminated 
signs, and high vacuum air pumps. Considerable is used in 
preparing certain explosives, e.g., mercury fulminate, which 
is used in percussion caps and detonators to explode powder 
and nitroglycerin. 

The use of mercury in thermometers depends on several 
facts, e.g., it is heavy, bright liquid and has a uniform 
change of volume with a change of temperature (between 
a wide range of temperature). The curve showing the rela¬ 
tion of volume and temperature is almost a straight line 
(Fig. 214), that is, the expansion of mercury is regular. 


364 


A BRIEF COURSE IN CHEMISTRY 


479. Two series of mercury compounds. — Mercury, like 
copper and iron, forms two classes of compounds — mer¬ 
curous and mercuric. 
The valence of mer¬ 
cury is + 1 in mer¬ 
curous compounds 
and + 2 in mercuric. 
Solutions of mercu¬ 
rous salts contain mer¬ 
curous ions (Hg+), 
and of mercuric salts, 
mercuric ion (Hg++). 

480. Mercurous 
and mercuric chlo¬ 
rides. — The mercu¬ 
rous salt (HgCl) is a 
white, tasteless 
powder, insoluble in water. It is used as a medicine under 
the name of calomel. It is formed as a white precipitate 
when a chloride and mercurous nitrate interact — a test for 
mercury in mercurous ions (Hg+). The equation is: — 



Temperature 

Fig. 214. — Curve showing regular change 


in volume of 
temperature. 


mercury with change of 


NaCl + HgNOs = HgCl + NaNOs 

This test is confirmed by adding ammonium hydroxide which 
blackens the precipitate. 

The mercuric salt (HgCh) is a white, soluble, crystalline 
solid. It is manufactured by heating a mixture of mercuric 
sulfate and sodium chloride, thus : — 


HgS 04 + 2 NaCl = HgCh + Na 2 S 04 

Mercuric Sulfate Sodium Chloride Mercuric Chloride Sodium Sulfate 


Mercuric chloride is a violent poison. The best antidote 
is the white of a raw egg. The albumin forms an insoluble 
mass with the poison. 

The common name of mercuric chloride is corrosive sub¬ 
limate (or bichloride of mercury). It has powerful antiseptic 
properties, and a dilute solution (1 gm. in 1000 gm. of water) 
is extensively used in surgery to sterilize instruments and to 
protect wounds from infection. 
















MAGNESIUM — ZINC — MERCURY 


365 


Mercuric chloride, when treated carefully with stannous 
chloride, is reduced first to white mercurous chloride and 
finally to a dark gray precipitate of finely divided mercury — 
the test for mercuric ions (Hg++). The equations for these 
reactions are: — 

2 HgCh + SnCh = 2 HgCl + SnCU 
2 HgCl + SnCh = 2 Hg + SnCh 


EXERCISES 

1 . State the properties and uses of magnesium. 

2 . Write equations for (a) interaction of magnesium and sulfuric 
acid and (6) heating magnesium in nitrogen. 

3 . Describe the metallurgy of zinc. 

4 . Summarize the (a) physical and (6) chemical properties of zinc. 

6 . What are the tests for zinc ? 

6 . What are the tests for (a) mercury, (6) mercurous compounds, 

(c) mercuric compounds. 

7 . Describe (a) mercurous chloride and (6) mercuric chloride. 
What is the commercial name of each? The use? 

8. What is (a) magnesia, (b) Epsom salts, (c) galvanized iron, 

(d) Chinese white, (e) white vitriol, (f) calomel, (ff) corrosive sublimate? 

9 . Write the ordinary and the ionic equations for (a) mercuric 
chloride and hydrogen sidfide form mercuric sulfide and hydrochloric 
acid, (b) magnesium chloride and sodium hydroxide form magnesium 
hydroxide and sodium chloride, (c) zinc hydroxide and sodium hydroxide 
form sodium zincate and water, (d) mercuric chloride and stannous 
chloride form mercurous chloride and stannic chloride. 

10 . Write the formulas of the following compounds by applying the 
principle of positive and negative valence, or by utilizing analogous for¬ 
mulas in this chapter: Magnesium bromide, magnesium nitrate, magne¬ 
sium sulfide, zinc chromate, zinc carbonate, zinc acetate, zinc phosphate 
(ortho), mercurous fluoride, mercuric sulfate, mercurous oxide. 


PROBLEMS 

1 . Calculate the per cent of the metallic element in (a) magnesium 
oxide, (b) zinc oxide, and (c) mercuric oxide. 

2 . Calculate the atomic weights of magnesium, mercury, and 
zinc : (a) 16.0263 gm. of MgO give 47.8015 gm. of MgS 04 ; (b) 16.03161 
gm. of zinc give 19.9568 gm. of ZnO; (c) 118.3938 gm. of HgO give 
109.6308 gm. of mercury. (Use exact atomic weights.) 


366 


A BRIEF COURSE IN CHEMISTRY 


SUGGESTIONS FOR LABORATORY WORK 

(References are to Newell’s Laboratory Exercises in Chemistry) 

Exercise 16 — Properties and Chemical Change. 

Exercise *21 — Equivalent Weight of Magnesium. 

Exercise S56 C — Test for Magnesium. 

Exercise S58c — Cobalt Nitrate Tests (Magnesium). 

Exercise S9c — Test for Zinc — T. 

Exercise 20 — Equivalent Weight of Zinc. 

Exercise S33c — Testing for Ions (Zinc). 

Exercise S34 — Hydrolysis of Certain Salts (Zinc). 

Exercise *61 c — Displacement of Metals (Zinc). 

Exercise 62d — Flame Tests for Metals (Zinc). 

Exercise S56 D — Tests for Metals (Zinc). 

Exercise S586 — Cobalt Nitrate Tests (Zinc). 

Exercise SI — Decomposition of Mercuric Oxide. 

Exercise *6la — Displacement of Metals (Mercury). 

Exercise S56 E — Tests for Metals (Mercury). 

Exercise S62 — Qualitative Analysis. 


TOPIC XXIV: SILVER —PHOTOGRAPHY-GOLD 

481. Introduction. — Silver and gold are precious metals. 
They have been used for ages in the form of ornaments, 
costly vessels, and coins. The Latin names of these metals 
are argentum and aurum, from which the symbols Ag and Au 
are derived. 

SILVER 

482. Occurrence of silver. — The chief ore is silver sulfide 
(argentite, Ag2S), which is usually associated with lead sulfide 
and other complex sulfide ores. In fact, silver-bearing ores 
are the main source of silver. 

483. Metallurgy of silver. — 1. Ores containing free silver, or silver 
compounds that can be easily changed into silver, are treated by the 
amalgamation process. The powdered ore is first changed, if necessary, 
into silver chloride by roasting with sodium chloride. The silver is 
displaced from the chloride by agitation with water and iron (or an iron 
compound), and then is extracted by adding mercury, which forms an 
amalgam with the silver (§477). When the amalgam is heated, the mer¬ 
cury distils off and the silver remains behind. 

2. In the Parkes process sulfur, arsenic, and other impurities are 
removed by roasting. The final mixture of lead, silver, and gold is 


SILVER — PHOTOGRAPHY — GOLD 


367 


melted with about 1 per cent of zinc. As the mixture cools, an alloy 
of silver, gold, zinc, and a little lead rises to the surface, solidifies, and is 
skimmed off. \ The skimmings are heated in a retort to volatilize the 
zinc, and then in a shallow furnace (cupel furnace) to convert the lead 
into an oxide (PbO), which melts and runs off, leaving an alloy of silver 
and gold (§ 496 and Fig. 220). This latter process of removing lead is 
called cupellation. Silver and gold are separated by electrolysis 
(§ 496). 

3. In the cyanide process pulverized ore is mixed with dilute sodium 
cyanide solution. The silver forms silver cyanide (AgCN) which reacts 
with the excess of sodium cyanide to form soluble sodium silver cyanide 
(NaAg(CN) 2 ). Zinc is added to precipitate the silver (§ 462), thus: — 

2 NaAg(CN )2 + Zn = 2 Ag + Na2Zn(CN)4 

4. Silver can be separated from its alloy with gold by several pro¬ 
cesses. (a) In the older process the alloy is boiled with concentrated 
nitric (or sulfuric) acid; the silver forms a soluble silver salt (nitrate or 
sulfate) but the gold is not acted upon. The silver is precipitated from 
the diluted solution by metallic copper. (6) In the electrolytic process 
the anode is a silver-gold alloy rich in silver and the solution is silver 
nitrate in nitric acid; the silver ion migrates to the cathode, and is 
deposited, while the gold is left as a powder or skeleton at the anode 
(§ 496). 

484 . Properties of silver. — Silver is a lustrous, white 
metal. It is harder than gold, but softer than copper. It 
is ductile and malleable, and can be easily made into various 
shapes. Silver has a specific gravity of about 10.5, being 
heavier than copper, but lighter than lead. It melts at 960° C. 
Silver conducts electricity the best of all metals, but it is too 
expensive for general use. 

486 . Chemical properties of silver. — Silver does not tar¬ 
nish in air unless hydrogen sulfide is present, and then the 
familiar brown (or black) film of silver sulfide (Ag 2 S) is pro¬ 
duced. It also turns black when in contact with organic 
sulfur compounds, e.g., those in perspiration, vulcanized 
rubber, eggs, and mustard. The tarnishing of household 
silver is caused by hydrogen sulfide in illuminating gas or 
gas from burning coal. 

Certain metals, i.e., those higher in the displacement 
series (§ 462 ), precipitate silver from its solutions. 

Silver is only very slightly acted upon by hydrochloric acid, 
and not at all by molten sodium hydroxide, potassium hy- 


368 


A BRIEF COURSE IN CHEMISTRY 


droxide, or potassium nitrate. Nitric acid and hot concen¬ 
trated sulfuric acid change it into the nitrate (AgNOs) and 
sulfate (Ag 2 S 04 ) respectively. Sodium cyanide in the pres¬ 
ence of air and water changes it into sodium silver cyanide 
(NaAg(CN) 2 ). 

Solutions of simple silver salts contain silver ions (Ag+), 
whereas complex salts yield complex ions, e.g., the silver- 
cyanogen ion (Ag(CN) 2 “) and the silver-ammonia ion 
(Ag(NH3)2+). 

486. Cleaning silverware. — Tarnished silverware can be cleaned by 
rubbing off the film of sulfide with a very soft abrasive — “ silver polish.” 
It can be safely and quickly cleaned by an electrolytic process, which 
has largely replaced the old process. A piece of metalhc aluminum and 
the tarnished object are immersed in a hot solution of sodium bicar¬ 
bonate and sodium chloride, and kept in contact. The cleaning is soon 
accomplished. The object is removed, thoroughly washed in clean, hot 

water, and dried. The pro¬ 
portions for household use are 
a teaspoonful each of baking 
soda and common salt to a 
quart of water. The best re¬ 
sults are obtained when the 
solution is very hot and the 
two metals are in good con¬ 
tact. 

487. Uses of silver.— 

Silver is too soft for con¬ 
stant use, and is hard¬ 
ened by adding a small 
amount of copper. These 
alloys are used as coins 
and for jewelry. Silver 
coins of the United States 
contain 900 parts of silver 
to 100 parts of copper, 
and are called “ 900 fine.^^ British silver coins were formerly 
“ 925 fine.” This quality is called sterling silver, and from 
it much ornamental and useful silverware is made. Large 
quantities of silver are used to plate other metals and to make 
silver compounds, especially silver nitrate (Figs. 215, 217). 



CouTtesy Eastman Kodak Co. 


Fig. 215. — Bars of silver in a large pho¬ 
tographic plant ready for making into 
silver nitrate. 




SILVER — PHOTOGRAPHY — GOLD 


369 


488. Silver plating. — The object to be plated is cleaned, 
and made the cathode in a solution of potassium (or sodium) 
silver cyanide; the anode is a plate of pure silver (Fig. 216). 
When the electric current 
is passed, the silver dis¬ 
solves from the anode and 
deposits on the cathode. 

The deposit of silver is 
dull, and is brightened 
by rubbing. 

Mirrors and reflectors 
(especially for automobiles) are made by coating glass with 
silver. A mixture of silver nitrate and ammonium hydroxide 
is reduced to metallic silver, which sticks to the glass and is 
protected by varnish. 

489. Silver nitrate. —This is a white crystalline solid, made 
from silver and nitric acid (Fig. 217). The equation is: — 

3Ag + 4HNa3 = 3AgN03 + NO + 2 H 2 O 

Silver Nitric Acid Silver Nitrate Nitric Oxide Water 




(jounesy isasiman iLoaas, vo. 
Fig. 217. — Making silver nitrate. 
























370 


A BRIEF COURSE IN CHEMISTRY 


It is soluble in water. It turns dark in contact with organic 
matter owing to partial reduction to metallic silver. For 
this reason it blackens the skin; if applied long enough, it 
disintegrates the flesh, and is used by surgeons to cauterize 
sores or abnormal growths. 

490. Silver halides. — The chloride, bromide, and iodide 
are formed as curdy precipitates when the silver ion (Ag+) 
unites with the halide ion (Cl~, Br“, I“). The chloride 
(AgCl) is a white, curdy solid, which soon turns violet and 
finally brown in the light. It is converted by ammonium 
hydroxide into a soluble compound (Ag(NH 3 ) 2 Cl), from which 
silver chloride can be reprecipitated by adding an excess of 
nitric acid. 

The bromide (AgBr) and iodide (Agl) are analogous to 
silver chloride in their properties and methods of formation. 
The bromide is pale yellow, and the iodide is yellow. They 
are used in photography (§ 491). 

The formation and properties of silver chloride constitute 
the test for silver (compare § 136). 

491. Photography. — This operation is based mainly on the fact that 
silver salts, especially the bromide and iodide, darken when exposed to 
the Hght. There are four steps. 

1. Exposing. The photograph is taken on a glass plate, or a film, 
coated on one side with a thin layer of gelatin containing very fine grains 
of silver bromide. The plate or film is quickly exposed in the camera. 
The light that comes from the object passes through the lens, forms an 
image on the plate or film, and changes the silver salt in proportion to 
the intensity of the hght reflected from the object. 

2. Developing. The plate or film is immersed in a solution of a 
mild reducing agent, called a developer, e.g., hydroquinone, pyrogallic 
acid, or special mixtures. As the developer acts upon the silver salt 
on the plate or film, the image gradually appears. This is a deposit of 
finely divided silver which varies in thickness in proportion to the light 
that fell upon the plate or film, being thickest where the hght was most 
intense. 

3. Fixing. The image is fixed by dissolving away the unchanged 
silver salt with a solution of sodium thiosulfate (or “ hyposulfite ”)> and 
then washing the plate or film thoroughly in running water. On the 
fixed plate the dark parts of the object appear hght and the hght parts 
dark; the plate in this condition is called a negative (Fig. 218, left). 

4. Printing. The print, or photograph, is made on paper coated 
with a mixture much hke the one on the plate, though less sensitive to 


SILVER — PHOTOGRAPHY — GOLD 


371 


light. The paper is laid upon the negative and both are exposed to the 
Ught so the hght will pass through the negative first. Since the negative 
obstructs the hght in proportion to the thickness of the silver deposit, 
the dark and light parts are reversed and a positive is obtained, i.e., a 
photograph which has approximately the same shading as the object 
(Fig. 218, right). 

On some kinds of paper the image appears at once, but on others it 
must be developed and fixed. Subsequent treatment called toning 
produces special results. 



Fig. 218. — Daguerre — a pioneer in photography (negative (left) and 
positive (right)). 


GOLD 

492. History of gold. — Gold, like silver, is one of the old¬ 
est metals. For ages it has been the most highly prized of 
the metals and extensively used for personal adornment. 
The Latin name of gold, aurum, gives the symbol Au. 

For several centuries the mediaeval chemists, or alchemists, 
practiced alchemy, Le., they tried to produce gold by the 
transmutation of base or cheaper metals. 

493. Occurrence of gold. — Its native compounds are few 
and rare. It is never found pure, being alloyed with silver 
and occasionally with copper or iron. It is disseminated in 





372 


A BRIEF COURSE IN CHEMISTRY 


fine, almost invisible, particles among ores of other metals, 
though not so abundantly as silver. Much gold is found 
in veins of quartz, and in the sand formed by the disintegra¬ 
tion of gold-bearing rocks. 

494. Mining and metallurgy of gold. — Gold was formerly mined 
by washing the gold-bearing sand in large pans or cradles. Now the 
sand is scooped up by huge dredgers and washed by machines. In 
placer mining and hydraulic mining, streams of water wash away the 

gold-bearing earth into long 
troughs or sluiceways. The 
lighter particles are washed 
away and the heavy particles 
together with the gold sink 
to the bottom where they are 
caught on cross bars. From 
this mixture, gold and silver 
are extracted with mercury 
(§§ 462, 477). 

In vein or quartz minin g the 
lumps of gold-bearing quartz 
are crushed to a fine powder 
in stamp mills, i.e., in a row 
of huge iron mortars by hard 
pestles (Fig. 219). The moist¬ 
ened mass is floated over 
copper plates coated with 
mercury, which collects or 
dissolves about 50 per cent of 
the gold. The amalgam is 
heated, as in the metallurgy 
of silver, to remove the mer¬ 
cury, and the gold is extracted 
from the residue. The rest 
(“ tailings ”) of the gold which is left in the slime is often extracted by 
the cyanide process (see below). 

In the chlorination process the ore is treated with water containing 
chlorine or with bleaching powder and sulfuric acid; this operation forms 
a soluble gold chloride (AuCls), from which the gold is precipitated as a 
fine powder. 

In the cyanide process the ore, usually low grade, or the slime from a 
previous extraction, is mixed with a weak solution of sodium cyanide 
and exposed to the air; this operation changes the gold into a soluble 
cyanide, thus: — 

4 Au -h 8 NaCN + O 2 + 2 H 2 O = 4 NaAu(CN )2 + 4 NaOH 

Gold Sodium Oxygen Water Sodium Gold Sodium 

Cyanide Cyanide Hydroxide 



Fig. 219. — Stamp mills for crushing gold 
ore to a fine powder. Cut away on the 
right to show stamps (2 and 5 are 
lifted). Crushed ore passes through 
sieve on the left and the gold is caught 
by mercury on the plates (front). The 
rocky refuse is washed away. 




















































SILVER — PHOTOGRAPHY — GOLD 


373 


The gold is separated from this solution by electrolysis or by precipi¬ 
tation. 

496. Refining gold. — Refining is accomplished by elec¬ 
trolysis. In one method, which is used in the United States 
mints, the solution in the cell is a mixture of gold chloride 
and hydrochloric acid, the anode is an alloy of silver and 
gold — rich in gold, — and the cathode is a thin sheet of pure 
gold. Gold is deposited on the cathode, and the silver forms 
silver chloride around the anode (Fig. 220). 



Fig. 220. — Purefying gold in the assay laboratory in New York City. 


The purity of gold is expressed in carats. Pure gold is 
24 carats fine. An alloy, for example, containing 22 parts of 
gold and 2 parts of copper is 22 carat gold. 

496. Properties of gold. — Pure gold is a yellow metal but 
the color of commercial gold varies with the alloying metal. 
It is the most ductile and malleable of all the metals. The 
leaf into which it can be beaten is very thin. In some 
cases 110,000 leaves have a thickness of only 1 centimeter. 
Gold is one of the heaviest metals, its specific gravity being 
about 19.3. It forms alloys with many metals; most 
gold ” is an alloy — some is not gold at all! 

Air, oxygen, hydrogen sulfide, and most acids do not 
attack it, and for this reason it is sometimes called a noble 




374 


A BRIEF COURSE IN CHEMISTRY 


metal. It is changed into gold chloride (auric chloride, 
AuCls) by chlorine (§ 136). With sodium cyanide, as 
described in § 494 , it forms sodium gold (aurous) cyanide 
(NaAu(CN) 2 ). 

497 . Uses of gold. — Pure gold is too soft for use as 
jewelry or coins, and it is usually alloyed with copper or 
silver. Gold coins contain gold and copper (§ 458 ). The 
United States standard gold coins contaiil 90 per cent gold 
and 10 per cent copper. Gold leaf of various grades is used 
to ornament books and signs. Jewelers use gold for many 
purposes; such gold varies from 12 to 22 carats in purity, 
though 14 or 18 carat gold is commonly used. 

498 . Gold compounds. — Gold forms two series of com¬ 
pounds — aurous and auric, in which the gold has the valence 
of 1 and + 3 respectively. Auric chloride (AuCU) in 
dilute solution is reduced by stannous chloride solution to a 
beautiful purple precipitate; the latter is called “ purple of 
Cassius,’’ and is colloidal gold. Its formation is the test 
for gold. In gold plating, which is much the same as silver 
plating, the solution contains the potassium gold cyanides 
(KAu(CN )2 and KAu(CN)4). 

EXERCISES 

1. Prepare a summary of the metallurgy of (a) silver and (6) gold. 

2. Describe in order the steps in taking a photograph. 

3. How is gold purified? What is 14 carat gold? 

4. Prepare a summary of the properties of (a) silver and (6) gold. 

5. Complete and balance: Ag 2 S 04 + KBr = K 2 SO 4 H-. 

Write the final equation in ionic form. 

PROBLEMS 

1. Calculate the weight of (o) silver in 1 kg. of silver nitrate and 
(6) gold in 1 gm. of potassium auricyanide. 

2. How many grams of gold in a 14-carat ring which weighs 14 gm. ? 

3. (a) What weight of silver can be obtained from an American ten- 

cent piece which weighs 2.44 gm. ? (6) How much gold from an Ameri¬ 

can gold coin weighing 3.75 gm. ? 


ALUMINUM — CLAY AND CLAY PRODUCTS 375 


SUGGESTIONS FOR LABORATORY WORK 


(References are to Newell’s Laboratory Exercises in Chemistry) 


Exercises Sll II (6) (1) 
Exercise *25 (c) 
Exercise *26 
Exercise 22 II (d) 
Exercise 23 
Exercise S33 (a) 
Exercise S46 (c) 
Exercise S50 (d) 
Exercise S56 K 
Exercise S59 (h) 
Exercise S60 
Exercise S62 


Test for Chlorides. 

Test for a Chloride. 

Insoluble Chlorides (Silver part only). 
Test for Hydrogen Chloride. 

Aqua Regia — T. 

Testing for Ions. 

Test for a Bromide. 

Test for an Iodide. 

Test for Gold. 

Testing Salts. 

Silver Salts and Photography — T. 
Qualitative Analysis. 


TOPIC XXV: ALUMINUM —CLAY AND CLAY PRODUCTS 

499. Occurrence of Aluminum. — About 8 per cent of 
the earth’s crust is combined aluminum; in abundance it 
ranks first among the metals and third among the elements 
(§ 326). All important rocks (except limestone and sand¬ 
stone) and many rock-forming minerals are salts of silicic 
acid in which aluminum is one of the metallic components. 
Among such minerals are feldspar, mica, and hornblende, 
which are constituents of the common rocks called granite 
and gneiss. Clay, which is formed by weathering rocks, is 
mainly aluminum silicate. 

Corundum and emery are impure aluminum oxide (AI2O3). 
Bauxite is an aluminum hydroxide (H4AI2O5) ; it is often 
colored red by iron oxide. Cryolite is sodium aluminum 
fiuoride (NasAlFe). 

600. Manufacture of aluminum. — Aluminum is manu¬ 
factured by the electrolysis of aluminum oxide (AI2O3). 
The purified oxide, which is prepared from bauxite, is dis¬ 
solved in melted cryolite, and when the current passes, 
aluminum is deposited at the cathode. This process was 
discovered in 1886 by a young American chemist named 
Hall and perfected by him. 


376 


A BRIEF COURSE IN CHEMISTRY 


A sketch of the apparatus is shown in Fig. 221. An open iron vessel 
(C, C) lined with carbon is the cathode. Connection with the cathode 
is made at D. The anode consists of several graphite bars {A, A, etc.) 
attached to a copper rod, by which they can be lowered as the graph¬ 
ite is consumed (by the liberated oxygen). 

The bottom of the box is first covered with cryohte, the anode is 
lowered, and the box is then filled with cryohte (to which calcium fluo¬ 
ride is added to lower the fusion 
point). When the current is 
turned on an arc is formed, and 
the resistance generates enough 
heat to melt the cryolite. Pure, 
dry aluminum oxide is now added, 
which dissolves in the cryolite 
and dissociates somewhat as elec¬ 
trolytes do in water. The oxygen 
goes to the anode and unites with 
the graphite. The graphite an¬ 
odes are gradually lowered, and 
eventually have to be replaced. 

The aluminum sinks through 
the cryolite, and collects as a 
liquid at the bottom of the vessel. 
The process is continuous, fresh aluminum oxide being added and the 
molten aluminum being drawn off at intervals. The cryolite and calcium 
fluoride are not decomposed as long as aluminum oxide is present. The 
aluminum produced is 98-99.5 per cent pure. 

601. Physical properties of aluminum. — Aluminum is a 
lustrous white metal. It is the lightest of the common 
metals. Its specific gravity is only about 2.6 while that of 
iron is 7.8. It is ductile and malleable, and is extensively 
made into wire, rods, tubes, sheets, and foil (as thin as 
0.0005 inch = 0.00127 cm.). It is a good conductor of heat 
and electricity. Compared with most metals, aluminum is 
rather hard and strong, being exceeded in this property only 
by iron and copper. 

Aluminum melts at about 658° C. In the semi-molten 
state it can readily be extruded, i.e., forced or pressed out 
through an opening into desired shapes. Just below the 
melting point it is brittle and can be ground into powder. 
Aluminum can be cast and welded, though it cannot be turned 
well in a lathe, nor can it be readily soldered to produce a 
permanent joint. 




0 



1 4 

59 ^ 

foS Offl 

i 4 

W -H 


■ ■ A 

■^11 1 

■ maI 

IK 



Fig. 221. — Sketch of the apparatus 
for the manufacture of aluminum by 
the electrolysis of aluminum oxide. 


















ALUMINUM — CLAY AND CLAY PRODUCTS 377 


602. Chemical properties of aluminum. — Aluminum is 
very slightly tarnished by air, owing to the protecting film 
of oxide that forms on the surface. Hydrochloric acid 
reacts readily with it, thus : — 

2 A1 + 6 HCl = 2 AICI 3 + 3 H 2 

Aluminum Hydrochloric Acid Aluminum Chloride Hydrogen 

Under ordinary conditions nitric and dilute sulfuric acids 
do not affect it; concentrated sulfuric acid acts upon it, 
forming aluminum sulfate. It displaces many metals from 
their solutions, e.g., copper (§ 462). Sodium chloride inter¬ 
acts with it, if dilute acids are present. With sodium and 
potassium hydroxides it forms aluminates and hydrogen, 
thus: — 

2A1 + 6NaOH = 2 NasAlOs + 3 H 2 

Aluminum Sodium Hydroxide Sodium Aluminate Hydrogen 


603. Uses of aluminum. — The varied properties of 
aluminum, especially its strength, lightness, and durability, 
adapt it to numerous uses, e.g., parts of military outfits, caps 
for jars and bottles, surgical instruments, tubes, fittings of 
boats, automobiles, and airships (using about one third of 
the production),,parts of opera glasses and telescopes, frame¬ 
work of cameras, stock patterns for foundry work, hardware 
samples, scientific apparatus, and large vessels for chemical 
processes (Figs, 222, 223). Aluminum cooking utensils are 
extensively used. 

Its silvery appearance and durability lead to its extensive 
use as an ornamental metal, both in interior decorative work 
and as numerous small objects. Aluminum leaf is used to 
letter book covers and signs, and the foil is used as a wrapper 
for food, candy, and tobacco. The powder suspended in an 
adhesive liquid forms a paint for steam pipes, radiators, 
storage tanks, smokestacks, and other metal objects exposed 
to heat or the weather. 

Large quantities of aluminum are consumed in purifying 
steel, Le., removing air bubbles from the molten steel, thereby 
preventing the formation of small holes (“ blow holes ”) 
in the castings. 


378 


A BRIEF COURSE IN CHEMISTRY 



Fig. 222. — Aluminum tanks used in chemical plants. Each tank holds 
about 6000 gallons. 

604. Alloys of aluminum. — Aluminum forms useful 
alloys, those with copper being the most important. The 
proportions of aluminum vary widely (from 5 to 95 per cent), 
thus giving a great variety of alloys. Those containing from 
5 to 10 per cent of aluminum are called aluminum bronze. 
They have a golden luster, are readily melted, and are highly 
resistant. They have many uses. Those containing 90 



Fig. 223. — Large containers made of aluminum for shipping liquids ivhich 
attack other metals. 
















ALUMINUM — CLAY AND CLAY PRODUCTS 379 



to 95 per cent of aluminum are called duralumin and are 
the aluminum of the trade. They are silver-white. These 
alloys are used in making castings and household ware. The 
framework of airships is duralumin (Fig. 224). 

An alloy with magnesium, called magnalium, contains 
from 75 to 90 per cent of aluminum. It is hard, light, 
attractive, and durable, and is used as parts of chemical 
balances and scientific instruments. 


Fig. 224. — View of-part of the interior of the airship Los Angeles showing 
the framework made of duralumin. 

606. Thermit. — Aluminum is‘a powerful reducing agent. 
This property is utilized in the manufacture of certain metals 
and in welding. 

When a mixture of chromium oxide, for example, and pow¬ 
dered aluminum is ignited at one point, the reduction pro¬ 
ceeds rapidly throughout the mixture and the intense heat 
fuses the chromium, which can be removed from the crucible 
in a lump; the aluminum oxide rises to the top of the metal 
as a slag. The equation is: — 

fcOs + 2A1 = 2Cr + AI2O3 

Chromium Oxide Aluminum Chromium Aluminum Oxide 

Other metals, e.g., manganese, titanium, molybdenum, and 
tungsten, hitherto rare or expensive, are similarly prepared. 






380 A BRIEF COURSE IN CHEMISTRY 

If a mixture of ferric oxide (Fe 203 ) and powdered alumi¬ 
num is ignited by a special mixture, a violent reaction occurs. 
The oxide is reduced to iron, which is melted by the intense 
heat (3000° C.) and protected by the layer of molten alumi¬ 
num oxide. By using a special form of apparatus the molten 
iron can be conducted from the crucible into a mold around 
a joint or fracture (Fig. 225). This method is quick and 

effective and is widely used to 
weld steel rails and repair frac¬ 
tures in machinery. 

Mixtures of aluminum and ox¬ 
ides, used for this purpose, are 
called thermit, and the method 
is known as the alumino-thermic 
method. 

606. Aluminum oxide. — This 
is the only oxide of aluminum 
(AI2O3). It is often called alu¬ 
mina. Its native forms are co¬ 
rundum and emery. Both are 
very hard substances, pure co¬ 
rundum ranking next to diamond. 
The transparent varieties of co¬ 
rundum have long been prized as gems, among them being 
the sapphire and ruby. The alumina which is manufactured 
from bauxite is a white, powdery substance. 

Emery was formerly used as an abrasive. But it has been 
largely replaced by an artificial oxide called alundum. This 
product is manufactured by heating alumina in an electric 
furnace. On cooling, the mass forms a hard solid which is 
extensively used as an abrasive. Without abrasives like 
alundum and carborundum (§§ 51, 447) many modern indus¬ 
tries, e.g., automobile and airplane, would be well-nigh 
impossible. 

When alumina, or any other compound of aluminum, is 
heated on charcoal with a blowpipe, then cooled, moistened 
with cobaltous nitrate solution, and heated again, the mass 
turns a beautiful blue color. This is a test for aluminum 
(compare § 474). 



Fig. 225. — Sketch of a crucible 
and mold in position for weld¬ 
ing a steel rail with thermit. 















ALUMINUM — CLAY AND CLAY PRODUCTS 381 


607. Aluminum hydroxide. — This is a white, jellylike 
solid formed by adding ammonium hydroxide to a solution of 
an aluminum salt, thus: — 

AICI 3 + 3 NH 4 OH = A1(0H)3 + 3 NH 4 CI 

Aluminum Ammonium Aluminum Ammonium 

Chloride Hydroxide Hydroxide Chloride 

Aluminum hydroxide is an example of a gelatinous colloid 
or gel (§ 446). It is insoluble in water, but it dissolves in 
strong acids and strong bases (in excess), forming respec¬ 
tively aluminum salts and aluminates. 

Aluminum hydroxide has weak basic and acid properties 
(compare § 472). An equation illustrating the basic prop¬ 
erty is: — 

A 1 ( 0 H )3 -h 3HC1 = AICI 3 + 3 H 2 O 

Aluminum Hydroxide Hydrochloric Acid Aluminum Chloride Water 

One illustrating the acid property is: — 

A 1 ( 0 H )3 + 3 NaOH = Na 3 A 103 + 3 H 2 O 

Aluminum Hydroxide Sodium Hydroxide Sodium Aluminate Water 

608. Aluminum sulfate. — This is a white solid prepared 
from clay or bauxite by heating with sulfuric acid. The 
crystallized salt has the formula Al 2 (S 04)3 . 18 H 2 O. It is 
used as a mordant in dyeing and as an ingredient of the 
mixture (called size) put on paper to prevent ink from 
spreading. 

A solution of aluminum sulfate has an acid reaction on 
account of hydrolysis (§ 281); the equation for the hydroly¬ 
sis is: — 

Al2(S04)3 + 6 H 2 O = 2 A 1 ( 0 H )3 + 3 H 2 SO 4 

Aluminum Sulfate Water Aluminum Hydroxide Sulfuric Acid 

Practical application of this reaction is made in purifying 
water. Upon adding aluminum sulfate and lime to impure 
water, the gelatinous aluminum hydroxide that is precipi¬ 
tated slowly settles and carries with it suspended particles 
and bacteria (§ 81). 

609. Alums. — A mixture of concentrated solutions of 
aluminum sulfate and potassium sulfate deposits crystals of 


382 


A BRIEF COURSE IN CHEMISTRY 


potassium alum or simply alum (Fig. 226). Its formula is 
K 2 Al 2 (S 04)4 or K 2 SO 4 . Al 2 (S 04 ) 3 . It is the type of a class 
of similar salts called alums, which can be prepared by mixing 
sulfates of univalent and trivalent metals {e.g., K, Na, NH 4 , 
and Al, Cr, Fe). For example, chrome (or chromium) alum 
is K 2 SO 4 . Cr 2 (S 04 ) 3 . 

Alums are rather soluble in water, and their solutions have 
an acid reaction owing to hydrolysis (§ 281). They crystal¬ 
lize as octahedrons (Fig. 226) and contain twenty-four mole¬ 
cules of water of hydration. 

Alum (and sometimes aluminum sulfate) is an ingredient 
of alum baking powders (§ 283); the acid needed to liberate 
carbon dioxide is formed by the hy¬ 
drolysis of the alum. 

Alums are used in dyeing and 
printing cloth, in tanning and paper 
making, as a medicine, for harden¬ 
ing plaster, in fireproofing wood and 
cloth, and in preparing aluminum 
compounds. Aluminum sulfate is 
displacing alum for many purposes, 
especially the purification of water. 

610. Mordants. — Aluminum hydrox¬ 
ide is extensively used as a mordant in 
dyeing. Some fibers, e.g., cotton, do not 
take up, or adsorb, a dye readily. Hence 
to prevent the dye from being washed off, 
a substance called a mordant is added to fix the dye within the fiber. 
Gelatinous colloidal hydroxides of aluminum (§ 607), iron, tin, and 
chromium are used. 

The cloth to be dyed is first impregnated with an aluminum salt, such 
as aluminum acetate, and then exposed to steam. This operation 
changes the aluminum salt into aluminum hydroxide, which is precipi¬ 
tated in the fiber of the cloth. The mordanted cloth is next passed 
through a vat containing a solution of the dye, which is adsorbed by 
the aluminum hydroxide, forming a colored compound. The latter is 
relatively insoluble and cannot be easily washed from the cloth, i.e., it 
is a fast color. 

Dyes of this kind are called mordant or adjective dyes. They differ 
from (a) insoluble dyes hke indigo which are precipitated within the 
fibers, and (6) direct or substantive dyes which produce colors in the 
fabric without mordants. 



Fig. 226. — Alum crystals 
deposited from a concen¬ 
trated solution. 












ALUMINUM — CLAY AND CLAY PRODUCTS 383 


611. Clay. — This is a more or less impure aluminum 
silicate, formed by the slow decomposition of rocks containing 
aluminum compounds, especially the feldspars. Pure, typi¬ 
cal feldspar is potassium aluminum silicate (KAlSisOg). 
The products of its decomposition are chiefly an insoluble 
aluminum silicate and a soluble alkaline silicate. The latter 
is washed away. The pure aluminum silicate which remains 
is kaolin (H 4 AlSi 209 or H 2 Al 2 (Si 04 ) 2 . H 2 O). Usually kaolin 
is mixed with particles of mica and quartz, calcium and 
magnesium carbonates, and iron compounds — the last 
giving the colors. This mixture, which varies in com¬ 
position, is known as clay. 

Kaolin undergoes hydration when mixed with water and 
finally acquires marked colloidal properties. In this so- 
called plastic state, the colloidal suspension is rather perma¬ 
nent and can be molded into various objects which retain 
their shape when dry. When heated, it loses its water of 
hydration, and does not melt (except at a very high tempera¬ 
ture), but becomes a permanently hard mass without plas¬ 
ticity. These two properties (plasticity when wet and hard¬ 
ness when heated) are the foundations of the ceramic 
industries. 

For ages porcelain and china have been made by mixing kaohn 
(free from iron), fine sand, and powdered feldspar, shaping the mass 
into the desired form by molds or on a potter’s wheel, and then heating 
(“ firing ”) in a kiln to a high temperature (Fig. 227). The mass when 
cool is hard and translucent (if thin), but porous. It is called bisque. 
To make it impervious to gases and hquids, its surface must be glazed. 
This is done by dipping it into a creamlike mixture of feldspar and silica, 
similar to that used for making the porcelain but more easily fused, and 
then heating it again. The thin coating melts, runs over the surface, 
penetrates the porous mass and fuses with it to some extent, and when 
cold finally forms a shiny, impervious glaze. 

In making pottery the raw materials are less carefully selected and 
prepared, and not heated to such a high temperature. The best grades 
can hardly be distinguished from porcelain, but usually pottery is much 
heavier and thicker. 

If less pure, plastic clay is used and heated to a moderate temperature, 
the product is known as earthenware or stoneware. This is a large 
class and includes tiles, terracotta, jugs, and flowerpots. Tins ware is 
porous and is sometimes glazed by throwing salt into the kiln just before 
the operation is over. The salt forms a fusible sodium aluminum sihcate 


384 


A BRIEF COURSE IN CHEMISTRY 



Fig. 227. — Interior of a chemical porcelain factory. 


upon the surface. The special grades of stoneware for use in chemical 
plants are made by a more careful procedure. 

Clay products used for construction include bricks, drain pipe, etc. 
They are made from impure clay and heated just enough to harden the 
mixture. The product varies with the clay, but is often colored red 
owing to iron oxide formed from the iron compounds in the unburned 
clay. Buff bricks are made from clay containing a small proportion of 
iron. Fire bricks and other material designed to withstand high tem¬ 
peratures are made from clay containing considerable sihca. 

EXERCISES 

1 . Starting with aluminum how could you prepare in succession 
AlClg, A1(0H)3, NasAlOs, AlCb, Al(OH)3, AI2O3, Al? 

2 . Describe the manuifacture of aluminum. 

PROBLEMS 

1 . Calculate the weight of aluminum in (a) 20 gm. of aluminum oxide, 
(b) 34 gm. of aluminum hydroxide. 

2 . What weight of aluminum can be obtained from 100 kilograms of 
bauxite (93 per cent Al(OH) 3 )? 

SUGGESTIONS FOR LABORATORY WORK 
(References are to NewelFs Laboratory Exercises in Chemistry) 

Exercise S12 (c) — Purification of Water — T. 

Exercise *16 — Effect of Shape on the Solubility of a Solid — T. 





LEAD 


385 


Exercise SI8 — Equivalent Weight of Aluminum. 

Exercise 56 (g) — Testing Baking Powders. 

Exercise S56 A — Tests for Metals (Aluminum). 

Exercise S58 (a) — Cobalt Nitrate Tests (Aluminum). 

Exercise S61 —Preparation and Properties of Aluminum Hy¬ 

droxide. 


TOPIC XXVI: LEAD 

612. History of lead. — Lead and its compounds have 
been used since the dawn of history. The Chinese have 
used it for ages to line chests in which tea is stored and trans¬ 
ported. The Romans, who obtained it from Spain, called it 
plumbum nigrum, i.e., black lead, and used it for conveying 
water just as we do today. The symbol Pb comes from the 
Latin word plumbum. 

613. Metallurgy of lead. — Lead is obtained from galena (PbS) — 
the most abundant ore — by first roasting the ore to change part of the 
sulfide to the oxide and the sulfate. Then the air is excluded and 
the temperature raised; the mixture reacts thus: — 

2 PbS + PbS04 + 2 PbO = 5 Pb + 3 SO 2 

Lead Sulfide Lead Sulfate Lead Oxide Lead Sulfur Dioxide 

Lead produced by this process is impure and must be refined. The 
impurities make the lead hard and unfitted for most uses. There are 
two processes of refining. 

1. The Parkes process is described in § 483, 2. The refined lead, left 
in the kettles, is cast into molds by machinery. 

2. In the electrolytic process the cathode is a sheet of pure lead, and 
the electrolytic solution is a mixture of lead fluosilicate (PbSiFe) mixed 
with gelatin. The anode is a slab of the impure lead. When the cur¬ 
rent passes, pure lead is deposited on the cathode and the other metals 
remain attached to the remnant of the anode or sink to the bottom. 

614. Physical properties of lead. — Lead is a blue-gray 
metal. When scraped or cut, it has a brilliant luster, which 
soon disappears, owing to the formation of a film of oxide or of 
basic carbonate. It is so soft it can be scratched with the 
finger nail and rubbed off as black particles. It discolors 
the hands, and when drawn across a rough surface it leaves 
a black mark. For this reason it is sometimes erroneously 
called black lead (§ 40). 


386 


A BRIEF COURSE IN CHEMISTRY 


Lead is not tough nor very ductile. But it is malleable, 
and can be rolled into sheets and pressed (while just below 
the melting point) into pipe. It is a heavy metal, its specific 
gravity being 11.4; with the exception of mercury, it is the 
heaviest of the familiar metals. It melts at 327° C. 

616. Chemical properties of lead. — Lead when heated 
strongly in air changes into oxides (mainly litharge, lead 
monoxide, PbO). Hydrochloric and sulfuric acids have 
little effect upon compact lead. Nitric acid produces lead 
nitrate (Pb(N 03 ) 2 ). 

In the presence of air, weak acids like acetic acid (or vine¬ 
gar) and acids from fruits and vegetables change it into 
soluble, poisonous compounds. Soft water containing air, 
carbon dioxide, ammonia, nitrates, or chlorides, dissolves 
lead; and lead pipes should not be used to convey rain water 
or water containing ground gases. Permanently hard water 
(§ 298) usually forms a hard coating and thus prevents 
further action. All lead salts are poisonous, and ultimately 
cause serious illness. 

Certain metals, e.gr., zinc and iron, precipitate lead from 
its solutions as a grayish mass, which often has a beautiful 
treelike appearance (§ 462). 

616. Uses of lead. — Lead is extensively used as pipe. 
Lead pipe is not only used to convey water to and from parts 
of buildings, but also as a sheath for electric wires, both 
overhead and underground. Sheet lead is used to cover 
roofs and to line sinks, cisterns, and the cells employed in 
some electrolytic processes. The lead chambers and some 
evaporating pans used in manufacturing sulfuric acid are 
made of sheet lead. Shot and bullets are lead (alloyed with 
a little arsenic to make them hard). Spongy lead is used in 
preparing the plates of storage batteries. Lead wool (very 
fine wire) is used to calk pipe joints. 

Lead is used to make many useful alloys, which as a rule 
are harder than lead. Those containing considerable lead 
are type metal (lead, tin, antimony), solder (lead and tin), 
Britannia and Babbitt metals, and pewter (§ 416). Solder 
melts at a lower temperature than lead and can be poured 
around a lead joint. Fusible metals contain lead (§417). 


LEAD 


387 


617. Lead oxides. — The monoxide (PbO) is a yellowish 
powder known as massicot, or a buff-colored crystalline mass 
called litharge. It is formed by heating lead in a current of 
air. It is made this way, though considerable is obtained 
as a by-product in separating silver from lead (§ 613) . Large 
quantities are used in manufacturing some oils and varnishes, 
flint glass, a glaze for pottery, and as the source of many 
lead compounds. 

The tetroxide (red lead or minium, Pb 304 ) is a red powder, 
varying somewhat in color and composition. The pure 
compound has the composition represented by the formula 
Pb 2 Pb 04 , Le., lead plumbate. It is prepared by heating 
lead or lead monoxide at the right temperature (about 450° C., 
but not over 545° C.). It is used in making flint glass. 
Pure grades are made into artists^ paint, but the ordinary 
grade, if mixed with the proper oil, is used to paint structural 
iron work (bridges, fences, gasometers, etc.), hulls of vessels, 
and agricultural implements. A mixture of linseed oil and 
red lead is used by plumbers and gas fitters to make joints 
tight. Orange mineral has about the same composition as 
red lead, though its color is lighter; its uses are the same. 

The dioxide (Pb02) is a brown powder formed by treating 
lead tetroxide mth nitric acid or by the action of bleaching 
powder on sodium plumbite (Na 2 Pb 02 , formed by dissolving 
lead hydroxide in sodium hydroxide). It is a strong oxidizing 
agent. Its formula is 0:=Pb=0. It is extensively used 
as the essential ingredient of the positive plate of electric 
storage batteries. 

618. Lead carbonate. — This is obtained as a white 
powder by adding sodium bicarbonate solution to a solu¬ 
tion of a lead salt. Sodium and potassium carbonates, 
however, produce a basic lead carbonate which has the 
composition corresponding to the formula Pb 3 ( 0 H) 2 (C 03)2 or 
2 PbC 03 . Pb(OH) 2 ). It is known as white lead and is used 
extensively in the paint industry. 

619. Paints. — A paint consists essentially of three ingredients: 
(a) a powder, (6) an oil, and (c) a filler. The powder, which is often 
called a pigment, is suspended in the oil, usually linseed oil, which 
hardens or “ dries ” (really oxidizes) to a tough film and sticks to a 


388 


A BRIEF COURSE IN CHEMISTRY 


surface. The powder, which is usually white lead (§ 620) with or with¬ 
out some colored substance, gives opacity, body, and color to the paint. 
It also fills the minute holes in the dried oil film and thereby assists in 
protecting the surface of the painted object from the action of oxygen 
and moisture. The filler is a white inert substance, e.g., barium sulfate, 
or kaohn, added to fiU in the pores and make the paint “ go farther.” 
Paint is often thinned with turpentine or oil just before using. 

620. White lead. — This is a heavy, white powder which 
mixes well with linseed oil. It is the basis of many colored 
paints, pigments being added to give the desired color. 
White lead paint has a marked covering power, i.e., it covers 
a surface well; it also dries to a good finish. But it darkens 
on exposure to hydrogen sulfide (which is often present in 
the air of cities) owing to the formation of black lead sulfide 
(PbS). In recent years other paint bodies, as the solids are 
called, have been mixed with, or substituted for, white lead, 
e.g.j zinc oxide (§ 471), kaolin, barium sulfate, and lithopone 
(a mixture of zinc sulfide and barium sulfate). These are 

white solids which do not darken 
in the air, and they often improve 
the paint in other ways, e.g., give 
it an impervious surface. 

621. Manufacture of white lead. — 
White lead is manufactured by several 
processes. 

1. In the Dutch process, perforated 
disks of lead, called buckles, are put in 
earthenware pots, which have a sepa¬ 
rate compartment at the bottom, con¬ 
taining dilute acetic acid (Fig. 228). 
These pots are arranged in tiers in a 
large building; spent tan bark is placed 
between each tier. The building is 
now closed except openings for the 
entrance and exit of air and steam. The fermentation of the tan bark 
produces carbon dioxide and moisture; heat is also Hberated. The heat 
volatiUzes the acetic acid, which changes the lead into lead acetate. 
The moist carbon dioxide converts the lead acetate into basic lead car¬ 
bonate or white lead. 

2. In another process, melted lead is blown (“ atomized ”) into a 
very fine powder by a jet of steam, and the powder is beaten for several 
days with acetic acid, carbon dioxide, and air. 




Fig. 228. — Earthenware vessel 
containing lead buckles to be 
made into white lead. Buckle 
before {lower) and after {upper) 
corrosion. 

















LEAD 


389 


3. In a third process, white lead is precipitated by passing carbon 
dioxide into basic lead acetate solution. 

622. Other lead compounds. — Native lead sulfide (PbS) 
is the mineral galena, the chief ore of lead. It resembles 
lead in appearance, but is harder and is usually crystallized 
as cubes, octahedrons, or their combinations (Fig. 229). It 
is obtained as a black precipitate by the interaction of hydro¬ 
gen sulfide (or other soluble sulfides) and a solution of a lead 
salt. Its formation is a test for lead. The chloride (PbCh) 
is a white solid formed by adding hydrochloric acid, or a 
.soluble chloride, to a cold solution of a lead salt. It dissolves 
in hot water (§ 136). The sulfate (PbS 04 ) is a white solid, 
formed by adding sulfuric acid, or a soluble sulfate, to a solu¬ 
tion of a lead salt. It is very slightly soluble in water, but 



Fig. 229. — Galena crystals (cube, octahedron and cube, octahedron). 

soluble in concentrated sulfuric acid, hence crude sulfuric 
acid sometimes contains lead sulfate. The white powder 
called sublimed white lead ” is about 75 per cent lead 
sulfate; it is used as a body for white paint. The chromate 
(PbCr 04 ) is a yellow solid formed by adding a solution of a 
lead compound to a solution of potassium chromate or potas¬ 
sium dichromate. It is sometimes called chrome yellow and 
is used as a pigment in the paint industry. Its formation 
serves as a test for lead. The arsenate (Pb 3 (As 04 ) 2 ) is a 
white solid, sparingly soluble in water; it is used as an in¬ 
secticide for leaf-eating insects. 

EXERCISES 

1. Describe the metallurgy of lead. 

2. Name three alloys which contain large proportions of lead. 
Name several alloys containing a minor proportion. 

3. State the tests for lead. 







390 


A BRIEF COURSE IN CHEMISTRY 


4 . Practical topics: (a) How would you test paint for lead? 
(6) What advantage has tin over lead for pipes? Lead over tin? 
(c) What is red lead? White lead? Black lead? (d) How would 
you show the presence or absence of lead in a lead pencil ? 


PROBLEMS 

1 . How many gm. of lead (a) in 200 gm. of galena, (6) in 1 kg. of 
litharge ? 

2 . A cube of lead is 6 cm. on each edge. How much does it weigh? 


SUGGESTIONS FOR LABORATORY WORK 


(References are to Newell’s Laboratory Exercises in Chemistry) 


Exercise S3 (a) 

Exercise 31 B (a), (b), (c) 
Exercise *47 (6) 

Exercise *48 II (6) 
Exercise S48 (6) 

Exercise *53 B (a) 
Exercise *61 (5),(c) 
Exercise S56 G 
Exercise S59 
Exercise S62 
Exercise S62 (c) 


Preparation of Oxygen from Lead Diox¬ 
ide. 

Insoluble Salts of Lead. 

Test for a Sulfide. 

Test for a Sulfide. 

■ Properties of Sulfides. 

• Reduction of Lead Oxide. 

Displacement of Metals. 

Tests for Metals (Lead). 

■Testing Salts (Lead part only). 
Qualitative Analysis. 

Tests for Lead. 


TOPIC XXVII: RADIUM — RADIOACTIVITY 

623. What is radium ? — Radium is a metallic element, 
which is a constituent of rare uranium-bearing minerals. 
The richest deposits are in the Belgian Congo district of 
Africa, which is now the source of supply. 

The proportion of radium in these minerals is only a few 
milligrams to the ton. But this small proportion is carefully 
extracted by a chemical process; the radium is then separated 
from the other metals as radium, chloride or bromide by 
fractional crystallization. 

The term radium as usually used means a compound, e.g., 
the commercial salt is radium bromide (RaBr 2 ). The price 
of radium compounds is high, about $75,000 a gram (of actual 
radium in the salt). 


RADIUM — RADIOACTIVITY 


391 


524. Discovery of radium. — About 1896, the French 
physicist Henri Becquerel discovered that uranium com¬ 
pounds affect a photographic plate in the dark or when 
wrapped in light-proof paper. Some minerals containing 
uranium compounds, particularly pitchblende, were later 
(1898) found by Madame Curie (Fig. 230) to be more active 
than uranium compounds. She studied pitchblende carefully, 
and subsequently in collaboration with her husband, Pierre 
Curie (1859-1906), extracted from this mineral a minute quan¬ 
tity of a new substance which was exceedingly active, many 
thousands of times more active than an equal weight of a 



Fig. 230. — Madame Curie (1867- ) in her laboratory. 


uranium salt. Madame Curie gave the name radium to its 
elementary constituent. 

tl26. General properties of radium and radium compounds. 

— Radium forms salts like those of calcium, e.g., a soluble 
chloride (RaCh) and hydroxide (Ra(OH) 2 ), and a relatively 
insoluble sulfate (RaS 04 ) and carbonate (RaCOs). Volatile 
radium compounds tinge the Bunsen flame red. 

626. Radioactivity. — Besides the properties just enu¬ 
merated, radium compounds have special properties which 
are conspicuously different from those of most substances. 
They are called radioactive properties. The term radio¬ 
activity is used to include these properties. 





392 


A BRIEF COURSE IN CHEMISTRY 



Fig. 231. — Dish of 
photographed (by 
room. 


radium bromide 
itself) in a dark 


1. Radium compounds spontaneously evolve considerable 
heat. A radium salt is always from 3 to 5° C. warmer than 

the surrounding air. It 
has been calculated that 
radium liberates enough 
heat every hour to raise 
a little more than its own 
weight of water from the 
freezing point to the boil¬ 
ing point. Moreover, 
this liberation of heat is 
kept up continuously. 

2. Radium compounds 
affect a photographic plate 
just as light does. If a 
tube containing a radium 
compound is left a short 
time on a photographic 
plate wrapped in black 
paper, or if a dish con¬ 
taining a radium compound is exposed to a plate in the dark, 
an image is produced when the plate 
is developed (Fig. 231). 

3. Radium compounds ionize the 
surrounding air, i.e., make it a con¬ 
ductor of electricity. For example, 
radium compounds discharge an elec¬ 
troscope. An electroscope (Fig. 232) 
contains two thin strips (D, E) of 
gold or aluminum, which separate 
when the electroscope is charged with 
electricity. Radium compounds, if 
brought near a charged electroscope, 
ionize the air, which passes in at A, 
affects the charging rod B, and passes 
out at C. Thus it permits the escape 
of electricity from the electroscope, 
i.e., discharges the electroscope; 
hence the leaves (E, D) fall together. 



Fig. 232. — Sketch of 
electroscope. 











RADIUM — RADIOACTIVITY 


393 


(The device G, F allows the electroscope to be charged or dis¬ 
charged.) The electroscope is used to detect radium com¬ 
pounds, and to determine the proportion of radium in mix¬ 
tures. 

4. Radium compounds make certain substances luminous. 
This is a conspicuous, and possibly the best-known property. 
Luminous watch, clock, and instrument dials, push buttons, 
door knobs, electric light chains, etc., consist essentially of a 
specially prepared crystalline zinc sulfide containing a minute 
quantity of a radium compound. (Sometimes a less expen¬ 
sive radioactive substance is used.) 

During the World War radium mixtures were indispensable 
in making luminous dials for airplanes, submarines, gun 
sights, signs, etc. Tubes and other vessels containing radium 
compounds glow in the dark. In museums the visitor some¬ 
times is guided by luminous arrows of radium paint through 
dark passageways into a small inclosure, where the radium 
compounds are seen glowing! 

627. Interpretation of radioactivity. — It was first thought 
that radioactivity was due to rays or radiations; hence the 
name radium. Many interesting experiments prove, how¬ 
ever, ( 1 ) that radioactivity is caused by the spontaneous 
emission of two kinds of electrically charged particles, called 
alpha (a) and beta (/3) particles, and ( 2 ) that the emission of 
beta particles is accompanied by pulsations in the ether 
called gamma ( 7 ) rays. 

628. Alpha particles. — The alpha particles are positively 
charged and are shot off from the radioactive substance with 
great velocity — in some cases as great as 14,000 miles a 
second. 

Each alpha particle carries two unit charges of positive 
electricity. Alpha particles are charged helium atoms (He++). 
If deprived of their charges, they become helium atoms and 
the weight of each is four times that of a hydrogen atom. To 
the alpha particles are due ffiany of the electrical phenomena 
of radioactive substances, e.g., ionizing air. 

Although alpha particles move very fast in straight lines, 
they do not travel very far. They are completely stopped 
by the time they have plowed through the air to a depth of 


394 


A BRIEF COURSE IN CHEMISTRY 


3 to 8 centimeters. They are almost entirely stopped by a 
thin sheet of paper and by aluminum leaf 0.1 mm. thick. 
The paths of alpha particles have been photographed as the 
particles shot through moist air (Fig. 233). 

529. Beta particles. — The beta particles are electrons, 
i.e., particles of negative electricity. They move in straight 
lines (at first) with great velocity, varying from 16,000 to 
180,000 miles a second — the maximum being nearly as 
great as the velocity of light (186,000 miles a second). The 
beta particles are very light, their weight being about yAt 
the weight of a hydrogen atom. The beta particles produce 
most of the photographic effect of radioactive substances. 
(Refer to §§ 357, 358, 424 for electrons.) 

530. Gamma rays. — The gamma rays have the least 
ionizing and photographic power, but they are the most 

penetrating. They 
pass readily through 
thick layers of metal, 
e.g., several inches of 
lead. Glass tubes 
containing radium 
Fig. 233. — Path of alpha particles in moist air. salts are inclosed in 

thick lead vessels to 
absorb the gamma rays. Gamma rays are not affected by 
a magnet; hence they have no electrical charges. 

531. Radium and energy. — The heat, electrical, and light 
energy associated with radioactive substances is due to re¬ 
lease of energy within the radium atom, resulting in an 
atomic disintegration and expulsion of fragments at a great 
velocity. Thus, the rapidly moving alpha and beta particles 
are suddenly stopped by the air, metals, and the radioactive 
solid itself, and heat and light result. 

532. Uses of radium. — The products from radium, espe¬ 
cially beta particles and gamma rays, have a powerful effect 
on living matter, and are utilized* to cure certain skin diseases 
and abnormal growths. The radium compound itself is so 
expensive it is seldom used directly. One of the products 
of the disintegration of radium is radon. And since radon 
gives off gamma rays constantly, radon is used instead of the 





RADIUM — RADIOACTIVITY 


395 


radium compound. The radon is pumped off at stated inter¬ 
vals into tiny glass tubes which can be placed upon or within 
the flesh. 

633. Radium is decomposing spontaneously. — Although 
radium is an element which possesses many properties like 
those of the other elements, it differs from most of the ele¬ 
ments in being unstable. That is, radium is slowly disinte¬ 
grating. One of the final products is the stable element 
lead. 

Other products of the disintegration of radium atoms are 
produced, as shown in § 636. One product, viz., radon (for¬ 
merly called radium emanation and niton), is given off con¬ 
tinuously by radium compounds. It condenses at the tem¬ 
perature of liquid air and glows in the dark; by raising the 
temperature, the gas can be recovered by blowing air through 
the tube. It belongs to the same group of elements as 
helium and argon, and has properties much like those of 
argon (§ 419). Radon is an unstable gas, and gives helium 
as one product. 

Since the two elementary gases, helium and radon, are 
produced spontaneously from all radium compounds, the 
radium atom itself must be disintegrating spontaneously. 
This means that the atom of one element (radium) is trans¬ 
forming itself into the atoms of two other elements (helium 
and radon). But unlike other chemical transformations, the 
decomposition of the radium atom cannot be hastened or 
retarded by the chemist. It goes on unceasingly and un¬ 
varyingly at temperatures between liquid air and the electric 
furnace. 

Furthermore, this transformation means that the element 
radium is slowly disappearing. The change is not rapid. 
Experiments show that half of a given weight of radium 
would disappear in about 1700 years, half of the residue in 
another 1700 years, and so on. Radon is also disintegrating. 
Its rate, however, is rapid; the half period, as it is called, 

is about 4 days. ^ . j 

634. The radium disintegration series. — As intimated 
above, helium and radon are not the only elements formed by 
the disintegration of radium. There are at least eight others, 


396 


A BRIEF COURSE IN CHEMISTRY 


all rather unstable and short-lived, except one (radium—G, 
or radio-lead). Radium itself is a product of the disinte¬ 
gration of uranium. Uranium, the heaviest of all the ele¬ 
ments, is regarded as the parent substance in the radium dis¬ 
integration series. The others in the series are given in 
§ 636. All are elements; some are very unstable, the half 
period being only a few minutes. Helium is not a member of 
the series, but is given off in some of the transitions. Elec¬ 
trons are also expelled in some cases. The final product of 
this disintegration series is radio-lead (§ 636). 

636. Some points of the theory of radioactive disintegra¬ 
tion. — We have already stated (a) that radium compounds 
are continuously giving off alpha and beta particles and 
gamma rays, (6) that the process of radioactive change is the 
result of a spontaneous disintegration of the atoms of radium, 
and (c) that radioactive changes are independent of external 
conditions {e.g., temperature) and the kind of chemical 
compound, and proceed at a definite rate. Hence it is 
concluded that radioactivity is a property of the atom 
itself. 

Radioactive changes involve not only the production of 
elements but the expulsion of alpha and beta particles, ^.e., 
the loss of a charged helium atom and an electron. An elec¬ 
tron has a negligible weight, but a helium atom has the weight 
4. Hence the loss of an electron leaves the weight of an 
atom unchanged, whereas the expulsion of a helium atom 
lowers an atomic weight by 4 units. 

636. Transformations in the uranium series. — The 
changes in this series are shown in the following scheme, 
which includes the symbol and atomic weight of each element, 
and the kind of particle emitted (a = alpha particle and = 
beta particle or electron). 


Ui -U- UXi -A UXs U 2 -V lo(nium) -A- 

238.2 234.2 234.2 234.2 230.2 

Ra A- Rn A- Ra • A A- Ra • B A Ra • C A 

226 222 218 i 214 214 

Ra • Cl —>- Ra • D —>- Ra • E —>- Ra • F —Ra • G 

214 210 210 210 206 


[Radio¬ 

lead] 


RADIUM — RADIOACTIVITY 


397 


Let us glance at two examples. 

1. In passing from uranium to radium, three helium atoms 
are lost. Therefore, the atomic weight of radium should 
be 12 units (3 X 4) less than uranium. From the scheme we 
see that uranium has the atomic weight 238.2 and radium 226. 

2 . Eight helium atoms are lost in passing from uranium 
to radium—G (the final product). Hence the final product 
in this series should have the atomic weight 206 {i.e., 238.2 
— (8 X 4). The atomic weight of radio-lead was found to 
be 206.08 by exceptionally accurate experiments. 

The atomic weight of lead from other {i.e,, non-radioactive) 
sources is 207.2, which is actually — not estimated — higher 
than that of radio-lead. 

The two forms of lead have identical chemical properties, 
e.g., form compounds exactly alike in exactly the same way. 
They differ only in atomic weight. Elements which differ 
in atomic weight but not in chemical properties are called 

isotopes or isotopic elements. 

EXERCISES 

1 . Prepare a summary of (a) radium, (6) radioactivity, and (c) ura¬ 
nium disintegration series. 

2 . State the properties of (a) alpha particles and (b) beta particles. 

3 . State the uses of radium compounds. 

PROBLEM 

1 . Calculate the weight of radium in 1 milhgram each of (a) radium 
bromide, (b) radium nitrate, (c) radium sulfate. 


APPENDIX 


1. The pressure of water vapor in millimeters of mercury 
is: — 


Tempera¬ 

ture 

Vapor 

Pressure 

Tempera¬ 

ture 

Vapor 

Pressure 

Tempera¬ 

ture 

Vapor 

Pressure 

12 

10.5 

17 

14.4 

22 

19.7 

12.5 

10.8 

17.5 

14.9 

22.5 

20.3 

13 

11.2 

18 

15.4 

23 

20.9 

13.5 

11.6 

18.5 

15.9 

23.5 

21.5 

14 

11.9 

19 

16.4 

24 

22.2 

14.5 

12.3 

19.5 

16.9 

24.5 

22.8 

15 

12.7 

20 

17.4 

25 

23.6 

15.5 

13.1 

20.5 

18.0 

25.5 

24.3 

16 

13.6 

21 

18.5 

26 

25.0 

16.5 

14.0 

21.5 

19.1 

26.5 

25.7 


The numbers in the Vapor Pressure columns are the values 
for a in formulas for the reduction of gas volumes, e.g., the for¬ 
mula in § 348, last paragraph, and in the following condensed 
formula: — 

_ V'jP' - «) 

760(1 + (0.00366 X t)) 

2. Books. — Starred (*) books are primarily for teachers, though 
many parts of these books are suitable for pupils. 

1 . *Alembic Club Reprints. University of Chicago Press. 
Nos. 2 (Dalton, Atomic Theory), 3 (Cavendish, Air), 4 (Avogadro, 
Molecules), 6 (Davy, Alkalies), 7( Priestley, Oxygen), 9 (Davy, 
Chlorine). 

2 . Chemical Discovery and Invention in the Twentieth Century. 
Tilden. E. P. Dutton & Co. 


398 















APPENDIX 


399 


3. Chemistry and Civilization. Cushman. E. P. Dutton & Co. 

4. Chemistry Applied to Home and Community. Beery. J. P. 
Lippincott Co. 

5. Chemistry in Agriculture. Chamberlain. The Chemical 
Foundation (N. Y.). 

6 . ^Chemistry in America. Smith. D. Appleton & Co. 

7. Chemistry in the Home. Howe and Turner. Charles 
Scribner’s Sons. 

8 . Chemistry in Industry. Vols. I and II. Howe. The 
Chemical Foundation (N. Y.). 

9. Chemistry in the Service of Man. Findlay. Longmans, 
Green & Co. 

10. Chemistry in the WorWs Work. Howe. D. Van Nostrand Co. 

11. Coal and the Coal Mines. Greene. Houghton, Mifflin Co. 

12 . ^College Chemistry. Newell. D. C. Heath & Co. Ad¬ 
vanced book for reference. 

13. Discoveries and Inventions of the Twentieth Century. Cressy. 
E. P. Dutton & Co. 

14. *Elements of Industrial Chemistry. Rogers. D. Van Nos¬ 
trand Co. 

15. Essays in Historical Chemistry. Thorpe. Macmillan Co. 

16. * Experiments in College Chemistry. Newell. D. C. Heath 
& Co. (Experiments based on No. 12.) 

17. Famous Chemists. Roberts. Macmillan Co. 

18. Famous Chemists. Tilden. E. P. Dutton & Co. 

19. Feeding the Family. Rose. MacmiUan Co. 

20. Food Products. Sherman. Harper & Bros. 

21. * Handbook of Chemistry and Physics. Chemical Rubber 
Co., Cleveland, Ohio. (Tables of physical and chemical data.) 

22 . How to Make Good Pictures. Eastman Kodak Co., Rochester, 
N. Y. 

23. Industrial Chemistry. Benson. Macmillan Co. 

24. Modern Chemistry and Its Wonders. Martin. D. Van 
Nostrand Co. 

25. * Nature of Matter and Electricity. Comstock and Troland. 
D. Van Nostrand Co. 

26. *New Era of Chemistry. Jones. D. Van Nostrand Co. 

27. Non-technical Chats on Iron and Steel. Spring. F. H. 
Stokes Co. 

28. ^Outlines of Industrial Chemistry. Thorp and Lewis. Mac¬ 
millan Co. 

29. Romance of the Atom. Harrow. Boni and Liveright. 

30. Romance of Chemistry. Foster. Century Co. 


400 


APPENDIX 


31. Romance of Modern Chemistry. Philip. J. P. Lippincott Co. 

32. Science Remaking the World. Caldwell and Slosson. Double¬ 
day, Doran & Co. 

33. Short History of Chemistry. Venable. D. C. Heath & Co. 

34. Source, Chemistry, and Use of Food Products. Bailey. P. 
Blakiston’s Son & Co. 

35. Story of Chemistry {The). Darrow. Bobbs-Merrill Co. 

36. Story of Copper {The). Davis. Century Co. 

37. Story of Gold {The). Meade. D. Appleton & Co. 

38. Story of Iron {The). Surface. D. Appleton & Co. 

39. Story of Iron and Steel {The). Smith. D. Appleton & Co. 

40. Story of Oil {The). Tower. D. Appleton & Co. 

41. Story of a Piece of Coal. Martin. D. Appleton & Co. 

42. *Study of Chemical Composition. Freund. Macmillan Co. 

43. Triumphs and Wonders of Modern Chemistry. Martin. 
D. Van Nostrand Co. 

44. What We Eat. Hawk. Harper & Bros. 

45. Wonder Book of Chemistry. Fabre. Century Co. 


INDEX 


Abrasives, 380 

Absolute alcohol, 199; temperature, 
266; zero, 266 
Acetates, 198 

Acetic acid, 198; test, 198 
Acetylene, 36, 56, 190; torch, 25, 26, 
303 ; welding, 25, 26, 304 
Acid, 123-129, 131-141; finding 

strength, 125; ions, 133; naming, 
127; reaction, 123; salt, 106, 129, 
210, 217, 309; strong, 141; test, 
123 ; typical, 108 ; weak, 141 
Adequate diet, 315 
Adjective dyes, 382 
Adsorption, 35, 382 
Air, 113-122; argon in, 116, 118; 
composition, 113, 116, 261; carbon 
dioxide in, 118; liquid, 19, 119- 
122; mixture, 116; oxygen in, 17, 
19, 116; water in, 117; weight of 
liter, 114 

Airplane, 183, 308, 309, 321 
Airship, 56, 120, 379 
Alabaster, 225 
Alchemy, 371 

Alcohol, 199; absolute, 199; de¬ 
natured, 199; ethyl, 182, 199; 
grain, 182; methyl, 198; solid, 
182, 198; test, ethyl, 200; wood, 
198 

Alkaline reaction, 124 
Allotropy, 36, 165 

Alloys, 203, 321; aluminum, 378; 
antifriction, 322 ; antimony, 321 ; 
bismuth, 322 ; copper, 354 ; gold, 
371, 373 ; lead, 386 ; steel, 246 
Alpha particles, 393, 396; path, 394 
Alum, 210, 381, 382 
Alumina, 380 
Aluminates, 381 

Aluminum, 240, 243, 368, 375-384; 
acetate, 382; alloys, 378 ; bronze. 


378; chloride, 377; hydroxide, 
375, 381; in rocks, 375; in steel, 
377; oxide, 345, 375-376, 380; 
paint, 248, 377; sulfate, 381; 
test, 380 
Alundum, 380 

Amalgam, 203, 363, 366, 372 
Amalgamation, 366 
Amethyst, 339 

Ammonia, 144-150 ; adsorption, 35 ; 
liquid, 145, 194; manufacture, 

146-148 ; synthesis, 146-148; test, 
146 

Ammoniacal liquor, 145, 188, 189 
Ammonium, 150; bicarbonate, 208; 
chloride, 106, 108, 115, 150; di¬ 
chromate, 114; ferric citrate, 251; 
hydroxide, 144-150; ion, 140; 
nitrate, 296; nitrite, 114; salts, 
150; sulfate, 115, 151 
Amorphous, 165; carbon, 32; sulfur, 
165 

Anhydride, 66, 128, 170, 343 
Anhydrite, 225 

Anhydrous, 66 ; ammonia, 145 
Animal charcoal, 34 
Anions, 131-141 
Annealing glass, 348 
Anode, 137, 351, 352 
Anthracene, 190 
Anthracite coal, 32 
Antichlor, 105 
Antimony alloys, 321 
Aquadag, 31 

Aqua regia, 109-110, 156 
Argentite, 366 
Argon, 116, 118 
Arrhenius, 131 

Arsenic, 202, 360; compounds, 321; 

insecticides, 198; trioxide, 321 
Asbestos, 341 
Atmosphere, 113, 268 


401 



402 


INDEX 


Atomic hydrogen, 56; numbers, 
327-331; theory, 278; weights, 
80, 278-279, 281, 283, 292-295, 
324-327, 397 

Atoms, 12, 278-279; and symbol, 80 ; 
and ions, 132; in molecule, 286- 
287; nucleus, 279 
Auric chloride, 374 
Automobile, 36, 183, 203, 247, 308, 
361, 369, 380 

Avogadro, 285; law, 285-287 

Babbitt metal, 322 
Bakehte, 333 

Baking powder, 210; alum in, 382 
Balancing equations, 88 
Barite, 177 

Barium dioxide, 260; sulfate, 362; 
test, 230 

Barometer, 265, 268 
Barytes, 177 

Bases, 123-129, 131-141; ions, 133; 
naming, 128; strong, 141; weak, 
141 

Basic reaction, 124 
Baiixite, 375 
Benzene, 189, 190, 199 
Benzyl chloride, 105 
Berzelius, 79 
Bessemer steel, 239 
Beta particles, 393-396 
Bismuth alloys, 322 
Bisque, 383 
Bittern, 335 
Bituminous coal, 32 
Black lead, 31 

Blast furnace, 232—235, 349-351 
Bleaching, 103-105, 277; mixtures, 
212 ; powder, 61, 103; sulfur 

dioxide, 171-172 
Bleach liquor, 211 
Blooms, 237 

Blowpipe, 192; flame, 192; oxy- 
acetylene, 303 ; oxy-hydrogen, 56 
Blue print, 251; stone, 355; vitriol, 
177, 355 

Boiler scale, 225 
Boneblack, 34 
Bones, 227, 318 
Books, 398^00 

Borax, 212-213, 345; bead, 213 
Bordeaux mixture, 355 
Bornite, 348 
Boron oxide, 345 


Boyle, 268; law, 267-269, 272 
Brass, 354 
Bread, 197 
Bricks, 384 
Brimstone, 163 
Britannia metal, 322 
British thermal unit, 18C 
Bromides, 334-336, 356 
Bromine, 334-336; water, 335 
Bronze, 354, 378 
Brown, Robert, 73 
Brownian movement, 73 
B.t.u., 180 

Bunsen, 191; burner, 188, 190; 

flame, 191 
Burettes, 125 
Burning, 37, 45 
Butane, 187 

Cadmium, 360 
Calcite, 216 

Calcium acetate, 198; acid phos¬ 
phate, 210; acid sulfite, 309; 
arsenate, 321; bisulfite, 309; 
carbide, 36, 303; carbonate, 216- 
219; carbonate, acid, 217; cyan- 
amide, 115; fluoride, 333, 334; 
hydroxide, 221-222; nitrate, 115; 
oxide, 219-221; phosphate, 227- 
230, 318; silicate, 234, 345; sul¬ 
fate, 224-225; sulfite, 171-172; 
test, 230 

Calculations from equations, 89 
Calomel, 364 

Calorie, large, 213, 298; small, 180, 
262, 298 

Calorimeter, 298 

Candle, flame, 301-302; wax, 186 

Cannizzaro, 293 

Caramel, 194 

Carat, 30, 373, 374 

Carbohydrate, 310 

Carbolic acid, 190 

Carbon, 29-38; and energy, 37, 
263 ; and life, 29 ; allotropic, 36 ; 
amorphous, 32; atomic weight, 
293; disulfide, 36, 167; gas, 35; 
monoxide, 190, 199; test, 33; 

tetrachloride, 105 
Carbona, 105 

Carbonate, 30, 109; test, 219 
Carbon dioxide, 31, 38-43; and 
baking powder, 210; and life, 41; 
and plants, 263 ; fire extinguishers, 



INDEX 


403 


42; in air, 118; liquid, 40; solid, 
40; test, 38, 118; weight of liter, 
285 

Carbonic acid, 40 ; ions, 209 
Carborundum, 36, 343-344 
Carnelian, 339 
Casing head gasolene, 184 
Cast iron, 235 

Catalysis, 19, 23, 54, 57, 114, 147, 174, 
199 

Cathode, 137, 352 
Cations, 131-141, 203 
Caustic soda, 210 
Cellophane, 307 \ 

Celluloid, 309 ' 

Cellulose, 104, 172, 197, 307-310; 

acetate, 309 ; nitrates, 307-308 
Cement, 222-224 
Cementite, 235 

Centigrade scale, 266 ; thermometer, 
266 

Chalcocite, 160 
Chalcopyrite, 160, 348 
Chalk, 218-219 

Charcoal, 33, 105 ; animal, 34, 194 ; 

fuel value, 182; sugar, 194 
Charles law, 266-267, 272 
Cheese, 195 

Chemical change, 5-7, 13, 19, 21, 258 ; 
atoms, 279 ; electrons, 280 ; energy, 
258; ionic, 134 
Chemical equation, 85-92 
Chemical properties, 5, 6 
Chemical reaction, 85 
Chemist, work, 2, 243 
Chemistry, 1, 4 
Chile saltpeter, 213 
Chinaware, 383 
Chinese white, 361 
Chlorauric acid, 109 
Chlorides, 103, 109, 110, 124; in¬ 
soluble, 109; names, 110; of lime, 
103; test, 94, 95, 134 
Chlorination process, 372 
Chlorine, 99-106, 372 ; and bromides, 
334-335 ; and water, 61, 74 ; aqua 
regia, 109; bleaching, 103-105; 
liquid, 105; manufacture, 99, 
211-212; water, 74, 101; weight 
of liter, 101 
Chloroform, 199 
Chlorophyll, 42, 232 
Chlorpicrin, 105 
Chrome alum, 382 


Chromium oxide, 379; steel, 247 

Cinnabar, 160, 362 

Clay, 73, 341, 375, 381, 383-384 

Cleaning silverware, 368 

Clinker, 224 

Crisco, 57 

Crystallization, 69 

Crystals, 69 

Coal, 32, 179; beds, 262, 263; dis¬ 
tillation, 145; formation, 262; 
fuel value, 180; gas, 187-190; 
stove, 181; tar, 188-190 
Cobalt, alloy, 247; nitrate, 362, 380 
Coins, 354, 366, 368, 374 
Coke, 34, 189; fuel value, 182; 

petroleum, 187 
Cold storage plant, 149 
Collodion, 308 

Colloidal solution, 72; suspension, 
306; particles, 73 

Colloids, 73, 382; aluminum hy¬ 
droxide, 381; clay, 383 ; gold, 374 ; 
silica, 341; silicic acid, 342, 343; 
starch, 196 

Combination, 21-23, 54 
Combustion, 22, 23, 102, 121; La¬ 
voisier’s experiment, 261; prod¬ 
ucts, 38; spontaneous, 23 
Common salt, 207 

Composition, 80; constant, 10, 278, 
280; qualitative, 75 ; percentage, 
81; quantitative, 75 
Compounds, 8-13, 15; oxidation, 21 
Concrete, 223 
Condenser, 62 

Conservation of matter, 278, 280 
Constant composition, 10, 278, 280 
Converter, copper, 351; steel, 239 
Copper, 348-356 ; alloys, 354, 374 ; 
and nitric acid, 157; and sulfuric 
acid, 169; blister, 351; bromide, 
356; compounds, 352, 354, 356; 
electrolytic, 352; ferrocyanide, 
353 ; nitrate, 157, 356 ; ores, 348 ; 
oxide, 352, 356; sulfate, 66, 355; 
test, 353 

Copperas, 177, 251 
Coral, 218 

Correcting gas volumes, 267-269 
Corrosive sublimate, 364 
Corundum, 375, 380 
Cream of tartar, 210 
Creosote, 190 
Crucible steel, 244 



404 


INDEX 


Cryolite, 375 

Cupellation, 367 

Cupric salts, 354 

Cuprous oxide, 356 

Curie, Mme., 391 

Curie, Pierre, 391 

Cutting metals, 304 

Cyanide process, 361, 367, 372 

Daguerre, 371 
Dalton, 278 
Davy, 205, 207 

Decomposition, 19; double, 125 
Definite proportions, 278 
Defiagration, 156 

Deliquescence, 71, 207, 210, 275, 359 
Denatured alcohol, 199 
Desiccator, 71 
Developer, 370 
Dewar flask, 119 
Dialysis, 343 
Diamond, 30, 343 
Diastase, 195, 197, 199 
Diatoms, 339 
Di-chlor-e thy 1-sulfide, 105 
Diffusion, 272 
Digestion, 313 
Di-phenyl-chlor-arsine, 105 
Disinfectants, 172, 190 
Displacement, 52; metals, 353, 356 ; 
series, 52, 203, 356 

Dissociation, ions, 131-141; mole¬ 
cules, 131-141 

Distillation, 62; coal, 145; de¬ 
structive, 188, 198 ; dry, 33 ; frac¬ 
tional, 184, 187; water, 62-63; 
wood, 197 
Dextrose, 195 
Double decomposition, 125 
Dry distillation, 33 
Dumas, 76 
Duralumin, 203, 379 
Duriron, 249 
Dutch process, 388 
Dyad, 95 
Dyeing, 382 
Dyes, 104, 382 
Dynamite, 341 

Earthenware, 383 
Effervescence, 68 
Efflorescence, 66, 208, 213, 275 
Electrode, 31, 137 

Electrolysis, 95, 126, 134, 136-139, 


203, 360; aluminum oxide, 375, 
376; copper sulfate,. 139, 351; 
chlorine, 99; gold chloride, 373; 
lead, 385 ; silver — gold, 367; 
sodium chloride, 211-212; sodium 
hydroxide, 138, 205 ; water, 74, 139 
Electrolytes, 126, 134 
Electromotive series, 203 
Electrons, 250, 280, 330 ; and atoms, 
279; and chemical change, 280; 
and radium, 396; beta particles, 
384; loss and gain, 250, 280; 
theory, 280 
Electroscope, 392 

Elements, 8, 9, 11-14, 278; allo- 
tropic, 36, 165; arrangement, 

324-331; atomic numbers, 327- 
331; bivalent, 94 ; families, 325- 
326; groups, 325-326; in body, 
258; in earth’s crust, 257; inert, 
35, 114; in ocean, 257; isotopes, 
397; molecular formula, 290; 
periodic arrangement, 324-327; 
quadrivalent, 95; transformation, 
395 ; trivalent, 95 ; univalent, 94 ; 
valence, 94-98 
Emulsion, 72 

Energy, 258, 375, 380; and carbon, 
263; and hydrogen, 264; and 
oxidation, 261; and radium, 394 ; 
from food, 310, 311, 313, 315 
Enzymes, 197, 199 
Epsom salts, 177, 359 
Equation, 13, 85-92; calculations 
from, 89-92 ; gas, 290; molecular, 
290; thermo-, 180, 263, 264, 303; 
volumetric, 290 
Equilibrium, 153 
Equivalents, 282 
Equivalent weights, 282-283 
Esters, 306 
Etching glass, 333 
Ethane, 187 
Ethanol, 199 

Ethyl acetate, 198, 200, 306 ; alcohol, 
182 

Ethylene, 188, 190 
Eudiometer, 275-276 
Evaporation, 65 

Fahrenheit, 65 

Faraday, 137-138 

Fats, 57, 310; hydrogenation, 264 

Fehling’s test, 195, 196 





INDEX 


405 


Feldspar, 341 
Fermentation, 195-197 
Ferric compounds, 249-253, 321 
Ferricyanides, 251-253 
Ferro-manganese, 236, 239 
Ferro-silicon, 240, 243 
Ferrocyanides, 251-253 
Ferrous compounds, 249-253; sul¬ 
fate, 158, 251; sulfide, 165 
Fertilizer, 115, 151, 214, 228, 240; 

complete, 230 
Filtering water, 60 
Filter paper, 307 
Fire extinguisher, 105 
Fixation of nitrogen, 114, 115 
Flames, 299-304; acetylene, 303; 
blowpipe, 192; Bunsen, 191; 
candle, 301-302 ; illuminating gas, 
300; luminous, 301; » oxidizing, 
192, 213; oxy-acetylene, 56, 303; 
oxy-hydrogen, 56 ; reducing, 192 
Flashlight powder, 358 
Flotation, 350, 360 
Flour, 197 
Fluorine, 332-334 
Fluosilicic acid, 345 
Flux, 232 

Food, 310-317 ; water in, 59, 311, 312 
Formulas, 12, 15, 80; and ions, 132; 
and valence, 97; and molecular 
weight, 81; gas volumes, 398; 
graphic, 98 ; molecules, 287 ; sim¬ 
plest, 289 ; structural, 98 ; writing, 
97 

Fossil, 262 

Fourdrinier machine, 309, 310 
Frasch, 161 
Fructose, 195 
Fruit sugar, 195 

Fuels, 37, 179-189; B.t.u., 180- 

183, 188 ; composition, 179 ; foods, 
310, 313, 314; oil, 182-186, 241; 
value, 180, 298, 299, 313 
Fuming nitric acid, 298 
Fusible metals, 323 

Galena, 385, 389 
Galvanized iron, 249, 361 
Gamma rays, 393-394 
Gangue, 233 

Gas carbon, 35, 189 ; coal, 187-190 ; 
equation, 290 ; formulas, 88 ; illu¬ 
minating, 190 ; liquefaction, 121; 
mask, 101; mustard, 105; natural, 


187-188; poison, 105; producer, 
46, 299; sneeze, 105; solution, 
67; tear, 105 

Gases, 264-272; atoms in molecule, 
287 ; change of volume, 265-266 ; 
diffusion, 272; finding volume, 
265-269; formula for reduction, 
398 ; liquefaction, 272; measuring, 
264; structure, 271; weight of 
volume, 269 

Gasolene, 183-187; blending, 184; 
casing head, 184; hydrocarbons, 
184, 187 

Gay-Lussac, 277; law, 276-277, 284- 
285, 286 

Gel, 343, 381, 382 

Glass, 345-348 ; etching, 333-334 ; 
flint, 345, 387; kinds, 345-348; 
lehr, 347 ; pyrex, 345 ; water, 341 
Glauber’s salt, 177 
Glazing, 383 

Glucose, 195 ; test, 195, 356 
Glycerin, 305 

Gold, 73, 109-110, 352, 361, 367, 371- 
374; alloys, 374; chloride, 109, 
372,374; colloidal, 374 ; leaf, 373 ; 
refining, 373 ; specific gravity, 64 ; 
test, 374 
Gram, 259 

Gram-molecular volume, 287 
Gram-molecular weight, 287 
Grape sugar, 195 
Graphic formula, 98 
Graphite, 30-32, 344; crucibles, 244 
Gray cast iron, 235 
Green vitriol, 177, 251 
Guncotton, 156, 307 
Gypsum, 160, 177, 224-225 

Hsemoglobin, 232 
Hall, 375 

Halogens, 332, 337 
Hard water, 59, 225 
Helium, 11, 56, 119-126; alpha par¬ 
ticles, 393 ; and radium, 393, 395 ; 
charged atoms, 393 
Hematin, 311 
Hematite, 232, 243 
Heptane, 187 
Hexane, 184, 187 
Hofmann apparatus, 75 
Hydrate, 66, 275 
Hydrated lime, 221 
Hydraulic main, 188 



406 


INDEX 


Hydrocarbons, 29, 182, 184, 187, 189, 
190, 300 

Hydrochloric acid, 106-110; elec¬ 
trolysis, 137; test, 110 
Hydrobromic acid, 336 
Hydrofluoric acid, 332-334 
Hydrofluosilicic acid, 344 
Hydrogen, 49-57 ; and energy, 264 ; 
and metals, 52, 356 ; and methanol, 
199 ; and valence, 94 ; atomic, 56 ; 
bromide, 335; chloride, 103, 106- 
108; dioxide, 260, 277; displace¬ 
ment, 52, 356; flame, 53, 54; 
fluoride, 333; from acids, 49; 
from bases, 51; from water, 50: 
in acids, 123; in coal gas, 188, 
190; in illuminating gas, 190; 
peroxide, 260, 277; sulfide, 165- 
167 ; test, 55 ; weight of liter, 53 
Hydrogenation, 57, 264 
Hydrolysis, 209; alums, 382; alu¬ 
minum sulfate, 381; borax, 213 ; 
soap, 305; sodium carbonate, 209 ; 
starch, 197 
Hydrometer, 175 
Hydrosol, 343 
Hydrosulfuric acid, 166 
Hydroxides, 51, 109 
Hydroxyl, 124, 133 
Hypochlorous acid, 101, 103, 104 
Hypo, 178, 370 

Ice, 65; manufacture, 148 
Iceless refrigerator, 148, 172 
Illuminating gas, 190 
Indicators, 123 
Infusorial earth, 339, 341 
Ingots, 240, 242, 243 
Ink, 251 

Insecticides, 198, 321, 355 
Interaction, 85 
Invar, 246 
Invert sugar, 194 

Iodine, 214, 336-337; solution, 336- 
337; test, 196, 337; tincture, 337 
Iodides, 336-337 
Iodoform, 199, 337 
Ionic theory, 131, 209 
Ions, 131-141, 166, 205, 206, 212; 
and electrolysis, 136-139; and 
valence, 140; charges on, 140 ; 
color, 136 ; in solution, 139; iron, 
252; migration, 137; table, 140, 
141; taste, 136; water, 209 


Ionization, 131-141; degree, 141; 

table, 141 
Ionium, 396 

Iron, 232-253 ; alloys, 240, 243 ; and 
steam, 51; cast, 235; carbon in, 
35, 235; galvanized, 249, 361; 
hydroxides, 251; ions, 248, 249; 
ores, 232 ; ore, smelting, 232-235; 
oxides, 232; photomicrographs, 
237; pyrites, 160, 169, 251; 

Russia, 248; rust, 247; rustless, 
248; spiegel, 236; sulfate, 251; 
sulfide, 6, 13, 160, 165; tinned, 
248 ; test, 252 ; wrought, 236 
Isotopes, 397 

Kaolin, 283 

Kerosene, 183, 184, 199; flashing 
point, 186; hydrocarbons, 187 
Kiln, 220-222, 224 
Kilogram, 259 
Krypton, 118 

Lacquers, 308 
Lactic acid, 195 
Lactose, 195 
Lampblack, 35 

Lavoisier, 24, 25, 51; on combustion, 
261 

Law, 271; Avogadro, 285-287 ; Boyle, 
267-269, 272; Charles, 266-267, 
272; conservation of matter, 13, 
278, 280; constant composition, 
10, 278, 280 ; definite proportions, 
278, 280; Gay-Lussac, 276-277, 
284-285 ; new periodic, 331; old- 
periodic, 325 

Lead, 385-389 ; acetate, 198 ; alloys, 
386; arsenate, 321, 389; black, 
31, 385 ; carbonate, 387 ; chloride, 
110, 389; chromate, 389; dioxide, 
387; fluosilicate, 385; gamma 
rays, 394; metallurgy, 385 ; mon¬ 
oxide, 387; nitrate, 297, 386; 
oxides, 260, 385-387; pencil, 30, 
31, 104; plumbate, 387; radio-, 
395, 396, 397; red, 387; salts 
poisonous, 386 ; sulfate, 385, 389 ; 
sulfide, 385, 389; tetroxide, 387; 
test, 389 ; white, 387-388 
Leguminous plant, 115 
Lehr, 347 
Levulose, 195 
Liebig condenser, 62 




INDEX 


407 


Lighthouse, 301 
Lignin, 172 
Lignite, 32 

Lime, 219-221; caustic, 220; chlo¬ 
ride of, 103; hydrated, 221; milk 
of, 222; quick, 220; slaked, 219, 
220 ; sulfur, 160 ; superphosphate, 
228 

Limekiln, 220^222 
Limestone, 216 ; caves, 217 
Limewater, 222 
Linseed oil, 387 
Liquid air, 119-122 
Liter, 260 
Litharge, 386, 387 
Lithopone, 362, 388 
Litmus, 123 
Lubricant, 31 
Lubricating oil, 184, 186 
Luminous salts, 393 
Lye, 211 

Magnalium, 358, 379 
Magnesia, 359 ; milk of, 359 
Magnesium, 358-359 ; bromide, 334- 
335; chloride, 359; hard water, 
225-227 ; hydroxide, 359 ; nitride, 
114, 358; oxide, 359; sulfate, 359 
Malt, 195 

Maltose, 195-196, 197 
Manganese dioxide, 18, 100 
Manganese steel, 247 
Marble, 216, 221 
Massicot, 387 
Matches, 320-321 
Matte, 351 
Mendelejeff, 324 

Mercuric compounds, 363, 364; ox¬ 
ide, 8, 13, 17, 260, 261, 363 
Mercurous compounds, 363, 364 
Mercury, 120, 362-365, 372; chlo¬ 
ride, 110, 364; fulminate, 363; 
Lavoisier’s experiment, 261; ni¬ 
trates, 363 ; sulfide, 362 ; sulfate, 
363; test, 364, 365 
Metallic luster, 202 
Metallurgy, 203 ; copper, 348-351; 
gold, 372; iron, 232-235; lead, 
385; silver, 366 

Metals, 201-203 ; and water, 50, 51 ; 
displacing power, 203 ; properties, 
202 ; test, borax, 213 
Methanol, 182, 187-190, 198-199 
Methyl salicylate, 306 


Metric system, 7, 259-260 
Mexican onyx, 218 
Migration, ions, 137 
Milk, 195; of lime, 222; souring, 
195; sugar, 195 
Minium, 387 
Mineral, 202 ; water, 59 
Mixture, 9, 10 
Moissan, 332 
Molasses, 194, 199 
Mole, 287 

Molecular equation, 290 
Molecular formulas, 287, 289, 290 
Molecular weights, 81, 285; and 
formula, 81; molar method, 288 ; 
vapor density, 286 
Molecules, 9, 279; dissociation, 131- 
141; gas, 271-272; in solutions, 
73 

Molybdenum, 244, 247 
Monads, 95 
Monel metal, 354-355 
Mordant, 104, 381, 382 
Morley, 75, 76 
Mortar, 222 
Moseley, 330 
Muriatic acid, 106, 108 
Mustard gas, 105 

Naphthalene, 189 

Natural gas, 184, 187 ; fuel value, 188 
Negative, photographic, 370-371; va¬ 
lence, 95, 96 
Neon, 118 

Neutralization, 124--126, 133 
Neutral reaction, 123 
Newton’s metal, 323 
Nickel, catalyst, 57; steel, 246 
Niton, 395 
Nitrates, 152-158 
Nitric oxide, 152-158, 297 
Nitric acid, 152-158; and copper, 
297; and metals, 157; fuming, 
298 ; manufacture, 152-155 ; test, 
155, 158 
Nitrides, 114 
Nitrocellulose, 155 

Nitrogen, 113-117; and life, 113, 
115; dioxide, 153-157, 297; fixa¬ 
tion, 114, 115 ; liquid air, 113, 116 ; 
oxides, 152-158, 173, 296-297; 

test, 144; tetroxide, 298; weight 
of liter, 114 
Nitroglycerin, 155, 341 



INDEX 


408 

Nitro-hydrochloric acid, 110 
Nitrous oxide, 296 
Noble metal, 373-374 
Nodules, 115 
Non-electrolytes, 134 
Non-metals, 201 
Normal salt, 129 
Normal temperature, 265 
Normal pressure, 265 
Nucleus, 279 
Nutrients, 310 
Nutrition, 310-317 

Octane, 187 
Oildag, 31 
Olein, 264 
Onyx, 218, 339 
Ooze, 219 
Open-hearth furnace, 241-243 ; steel, 
240-244 

Orange mineral, 387 
Ore, 202 
Oxalic acid, 47 

Oxidation, 21-25, 54, 55, 155, 171; 
and energy, 261; broad meaning, 
249-250 

Oxides, 21, 108; and water, 128 
Oxidizing agent, 21 
Oxidizing flame, 192, 213 
Oxy-acetylene flame, 303-304 
Oxygen, 17-27 ; and digestion, 313 ; 
and life, 25-27; and plants, 263; 
atomic weight, 281; atoms in 
molecule, 286 ; from hypochlorous 
acid, 101; in air, 116-117, 261; 
ions, 206; molecular weight, 286 ; 
preparation, 17-19, 261; sources, 
260 ; test, 21; vapor density, 286 ; 
weight of liter, 19, 265, 269, 274, 
285 

Oxy-hydrogen flame, 56 

Paint, 362, 377, 387-389 ; aluminum, 
248; protective, 31, 248; radium, 
393; white, 362 

Paper, 309; board, 342; manu¬ 
facture, 172 

Paraffin wax, 183, 186, 187 
Paris green, 198, 321 
Parkes process, 366, 385 
Parting process, gold, 367 
Pentads, 95 
Pentane, 184, 187 
Percentage composition, 81-82 


Permanent hardness, 226, 359 

Permutit process, 226 

Periodic classification, 337 

Periodic law, 326, 331; new, 331 

Periodic table, 326 

Periods, 325 

Petrified wood, 339 

Petroleum, 182-187 ; cracking, 184 ; 

coke, 187 
Pewter, 386 
Phenol, 190 

Phenol-phthalein, 123-125 
Phlogiston, 24 
Phosgene, 105 
Phosphate fertilizer, 228 
Phosphate rock, 227-228, 318 
Phosphoric acid, 227, 230, 320 
Phosphorus, 116-117, 318-321; and 
air, 113 ; and life, 227 ; bomb, 320 ; 
cycle, 227; flame, 303; in steel, 
240; pentoxide, 319; red, 320 ; 
white, 319; sulfide, 320; yellow, 
319 

Photography, 178, 336, 370-371 
Photomicrographs, iron, 237 • steel, 
246 

Phylloxera, 160 
Physical properties, 4, 5 
Pickling iron, 251 
Picric acid, 156 
Pig iron, 235 
Pitchblende, 391 
Plaster of Paris, 225 
Plasticity, clay, 383 
Platinite, 247 

Platinum, 247 ; catalyst, 174 
Plumbago, 31 
Porcelain, 383 

Positive, photographic, 371; valence, 
95, 96 

Potassium, 50; acid fluoride, 332; 
acid tartrate, 210; alum, 382; 
bromide, 336; chlorate, 18, 81; 
ferricyanide, 251; ferrocyanide, 
251-253; fertilizer, 214; gold 
cyanide, 374; iodide, 337; salts, 
214 ; test, 50, 206 ; thiocyanate, 252 
Pottery, 383 
Powder, smokeless, 308 
Precipitate, 110 
Precipitation, 110 
Priestley, 17, 261 

Problems, weight, 89-90; volume, 
90-92 




INDEX 


409 


Producer gas, 46, 299 
Propane, 187 
Proper diet, 317 
Properties, 4, 5, 6 
Protein, 310, 314 
Prussian blue, 252 
Ptyalin, 196, 197 
Pulmotor, 26 
Purple of Cassius, 374 
Putty, 219 
Pyrene, 105 
Pyrex glass, 345 
Pyridine, 199 
Pyrites, 169, 251 
Pyroligneus acid, 198 

Quadrivalent element, 95 

Quartz, 338-341; crystals, 339; 

fused, 339-341 
Quicklime, 220 
Quicksilver, 363 

Radical, 80, 124, 132-141, 150; va¬ 
lence, 94-98 

Radioactive elements, 396 
Radioactivity, 391-397 
Radium, 11, 390-397; and atomic 
weights, 397; and energy, 394; 
bromide, 390, 392 ; decomposition, 
395-396 ; discovery, 391; emana¬ 
tion, 395 ; half period, 395 ; salts, 
391 

Radon, 394-396 
Rayon, 199, 307 
Reacting weights, 282 
Reaction, 13, 85; acid, 123; alka¬ 
line, 124 ; basic, 124 ; ionic, 134 ; 
neutral, 123; reversible, 153; 
velocity, 147, 174 

Reducing agent, 35, 370, 379; flame, 
192, 213; sugars, 195 
Reduction, 35, 55; broad meaning, 
249-250; by carbon, 35; by hy¬ 
drogen, 55 

Refrigeration, ammonia, 148; sulfur 
dioxide, 172 
Replacement, 52, 203 
Reverberatory furnace, 236, 350 
Reversible reaction, 153, 171 
Richards, 295 
Rose’s metal, 323 
Russia iron, 248 
Rusting, 20, 247-249 
Rustless iron, 248 


Saleratus, 210 
Sal soda, 208 
Salt, common, 207 

Salts, 50, 123-129, 131-141; acid, 
129, 210; formation, 128; ions, 
133 ; insoluble, 129 ; naming, 127 ; 
normal, 129; soluble, 128 
Saltpeter, Chile, 213 
Saponification, 305 
Scale, boiler, 225; thermometer, 266 
Selenite, 225 
Shellac, 199 

Silica, 338—341; colloidal, 341; gel, 
174 

Silicates, 341; sodium, 340-342 
Silicic acid, 73, 106, 341, 343-344; 
colloidal, 342 

Silicon, alloy, 249 ; carbide, 36, 343 ; 
dioxide, 338-341; test, 345; tet¬ 
rachloride, 106; tetrafluoride, 333, 
346 

Silver, 352, 366-371; bromide, 336, 
370 ; complex salts, 368; chloride, 
110, 295, 366, 370; cleaning, 368; 
iodide, 337, 370 ; nitrate, 368, 369 ; 
plating, 369 ; polish, 368 ; sterling, 
368; sulfate, 368; sulfide, 366, 
367; tarnishing, 367; test, 134, 
370 

Simplest formula, 82, 289 
Slag, 233, 234; copper, 350 
Slaked lime, 220 
Smithsonite, 360 
Smokeless powder, 308 
Smoke screen, 106, 344 
Sneeze gas, 105 

Soap, 57, 213, 225, 305, 306, 341 
Soda, 208-209, 210 ; ash, 209 ; bak¬ 
ing, 210; calcined, 209; caustic, 
210; cooking, 41, 210; crystals, 
208 ; lime, 106 ; lye, 211; sal, 208 ; 
silicate of, 340-342; washing, 41, 
208, 209 ; water, 39 
Sodium, 50, 205-214; acetate, 198; 
acid carbonate, 208-210; acid 
phosphate, 210 ; acid sulfite, 171; 
aluminate, 377, 381; and water, 
50, 206; bicarbonate, 41, 208- 
210; carbonate, 41, 207-209; 

carbonate, hydrolysis, 209; car¬ 
bonate, normal, 208 ; chloride, 207 ; 
chloride, electrolysis, 99, 211-212; 
cyanide, 372; gold cyanide, 372; 
hydroxide, 51, 205, 210-212; hy- 





410 


INDEX 


Sodium, {Continued) 

pochlorite, 104, 211; hyposulfite, 
178; iodate, 336; nitrate, 115, 
152, 213; nitrite, 114; peroxide, 
206, 260; plumbite, 387; silver 
cyanide, 367, 368; sulfite, 105, 
171; test, 50, 206; tetraborate, 
212-213; thiosulfate, 70, 178, 370 
Soft water, 59, 225 
Sol, 343 
Solder, 386 
Solid alcohol, 182 
Solids, solution, 68 
Solubility, 67-71; curve, 69, 70 
Soluble chloride, 110 
Solute, 67, 73 

Solution, 67 ; colloidal, 72, 73 ; con¬ 
centrated, 67; curve, 69, 70; di¬ 
lute, 67; ions in, 139; nature, 71- 
73 ; saturated, 67-70 ; supersatu¬ 
rated, 70; table, 68 
Solvay process, 207 
Solvent, 67 
Special steels, 246 
Specific gravity, 64; water, 64 
Spectrum, X-ray, 329-330 
Spelter, 360 
Sphalerite, 160, 360 
Spiegel iron, 236 
Spontaneous combustion, 23 
Spraying water, 61 
Sprinkler heads, 323 
Stainless steel, 247 
Stalactites, 218 
Stalagmites, 218 
Stamp mill, 372 
Standard atomic weight, 281 
Standard conditions, 265, 269 ; pres¬ 
sure, 269; temperature, 265 
Starch, 73, 196, 199; test, 196 
Steam, 65 
Stearin, 264 

Steel, 238-247 ; Bessemer, 239 ; chro¬ 
mium, 247 ; composition, 244-245 ; 
crucible, 244 ; electric process, 244 ; 
heat treatment, 245; high speed, 
247 ; manganese, 247 ; mild, 245 ; 
nickel, 246; open-hearth, 240- 
244; regenerative process, 243 ; 
soft, 245; special, 244, 246-247; 
stainless, 247; structural, 245; 
tempering, 245; tool, 245; tung¬ 
sten, 247 ; vanadium, 244, 247 
Stellite, 247 


Sterling silver, 368 
Stoneware, 383 
Stoves, blast furnace, 233 
Strontium, test, 230 
Structural formula, 98 
Stucco, 225 
Sublimation, 151, 336 
Substance, 4 
Substitution, 52, 133 
Sucrose, 194 

Sugar, 194-196; charcoal, 194; fer¬ 
mentable, 198 ; fruit, 195 ; grape, 
195 ; invert, 194 ; milk, 195 ; reduc¬ 
ing, 195; test, 195 

Sulfates, 177; acid, 177; normal, 
177; test, 135, 177-178 
Sulfides, 160, 166 

Sulfites, 170-172 ; acid, 171; normal, 
171 

Sulfur, 5, 6, 36, 160-165 ; and nitric 
acid, 157; atoms in, 163; chlo¬ 
ride, 163, 164 ; dioxide, 169-175 ; 
flowers, 163 ; in steel, 240 ; mining, 
160-163; monochloride, 163; 
monoclinic, 164; orthorhombic, 
164 ; plastic, 165 ; rhombic, 164 ; 
trioxide, 174 

Sulfuric acid, 172-177; and metals, 

176 ; contact process, 174 ; fuming, 
175 ; lead chamber, 173 ; test, 135, 

177 

Sulfurous acid, 170-172 
Superphosphate of lime, 228-230 
Suspension, 72 ; colloidal, 72 
Symbols, 12, 79; ions, 132; table. 
Cover 

Table, alloys, 354; atomic numbers, 
328; atomic weights. Cover; 
elements, 12, 257, 258; equivalent 
weights, 283; food, 311-317 ; ioni¬ 
zation, 141; ions, 140, 141; metals, 
non-metals, 201; periodic, 326; 
valence, 96 ; vapor pressure, 398 
Tar, coal, 188-190 
Tear gas, 105 
Temporary hardness, 226 
Test, acetic acid, 198; acid, 123; 
aluminum, 380; ammonia, 146; 
barium, 230; calcium, 230; car¬ 
bon, 33; carbon dioxide, 38, 118; 
copper, 353; ethyl alcohol, 200; 
ferrous compounds, 252; ferric 
compounds, 252; glucose, 356; 




INDEX 


411 


gold, 374; hydrochloric acid, 110; 
hydrogen, 55; hydrogen chloride, 
108 ; hydrogen sulfide, 167 ; iodine, 
337 ; ionic, 134 ; iron, 252; lead, 
389; mercury, 364, 365; metals, 
borax, 213; nitric acid, 155, 158; 
nitric oxide, 297; nitrogen, 144; 
nitrous oxide, 296; oxygen, 21 ; 
potassium, 50, 206; silicon, 345; 
silver, 370; sodium, 50, 206; 

soluble chloride, 110; strontium, 
230; sugar, 195; sulfate, 135; 
177-178; sulfuric acid, 135, 177 ; 
zinc, 362 
Tetrads, 95 

Theory, 271; atomic, 278 ; electron, 
280; ionic, 131, 209; kinetic, 
271; molecular, 271 
Thermit, 379-380 
Thermometer, mercury, 363 
Thermos bottle, 120 
Thomas-Gilchrist process, 240 ; slag, 
240 

Tin, block, 63 
TNT, 156 
Toluene, 189 
Toning, 371 
Travertine, 218 
Triads, 95 

Tri-nitrotoluene, 156 
Tripoli powder, 339 
Tungsten steel, 247 
Turnbull’s blue, 252 
Turpentine, 103 
Tuyeres, 233 
Type metal, 322, 386 

Ultramicroscope, 73 
Uranium, 391, 396-397; series, 

396-397 

Valence, 94-98, 249-250; and ions, 
140; and oxidation, 250; and 
reduction, 250; negative, 95, 96; 
positive, 95, 96; signs, 95-98; 
tables, 96 ; zero, 95 
Vanadium silicate, 174; steel, 244, 
247 

Vapor density, 286; pressure, 273- 
275, 398 

Vaseline, 183, 186, 187 
Velocity, reaction, 147, 174 
Vinegar, 198 
Viscose, 307 


Vitamins, 315 

Vitriol, blue, 177; green, 177, 251; 
white, 177 

Volumetric equation, 290 

Water, 18, 59-76; and fluorine, 333; 
and metals, 50, 51; and oxides, 
128; and sodium, 50, 206; and 
temperature, 64; as solvent, 67- 
73; boiling point, 63, 65, 274; 
carbonated, 40; chemical proper¬ 
ties, 66; chlorination, 61; chlo¬ 
rine, 74, 101; composition, 74-76, 
275-276; distilled, 62; electrol¬ 
ysis, 74, 139; freezing point, 63, 
65; gas, 299; glass, 341; hard, 
59, 359; hardness, 225-227, 359; 
hydrogen sulfide, 166 ; impure, 59 ; 
in body, 59, 310; in food, 59, 311, 
312; ions, 133, 209; lime, 222; 
maximum density, 64; metric 
weights, 260; mineral, 59; mole¬ 
cules, 73; of hydration, 66, 275, 
383; physical properties, 63; 
purification, 60-63, 251, 355, 381 ; 
rain, 59; royal, 109; salt, 60; 
soda, 39 ; soft, 59, 225 ; softening, 
225; sterilizing, 61; vapor, 65, 
117, 398; vapor pressure, 273- 
275, 398 

Weight of liter, air, 114; ammonia, 
145 ; carbon dioxide, 39 ; chlorine, 
101; hydrogen chloride, 108; ni¬ 
trogen, 114; oxygen, 19, 269; 
sulfur dioxide, 170 
Welding metals, 26, 56, 304 ; thermit, 
379 

Whey, 195 
White cast iron, 235 
White lead, 387-388; vitriol, 177, 
362 

Whitewash, 222 
Whiting, 219 
Wintergreen, 306 

Wood, alcohol, 198 ; distillation, 197; 
fuel value, 182 ; paper, 309 ; petri¬ 
fied, 339; pulp, 104 
Wood’s metal, 323 

World War, 105, 147, 320, 344, 393 
Writing equations, 86-89 
Wrought iron, 236 

Xenon, 118 

X-rays, spectra, 329-330 




412 


INDEX 


Zero, absolute, 266 
Zinc, 360-362 ; Amalgam, 363 ; am¬ 
monia-hydroxide, 362; blende, 
160, 360 ; chloride, 362 ; dust, 360 ; 
granulated, . 360 ; hydroxide, 362; 


oxide, 361; sulfate, 50, 362; sul¬ 
fide, 360, 362; test, 362; white, 
361 

Zincates, 361 
Zymase, 199 




I 


Deacidified using the Bookkeeper process. 
Neutralizing agent: Magnesium Oxide 
Treatment Date: June 2013 


PreservationTechnologies 

A WORLD LEADER IN COLLECTIONS PRESERVATION 



111 Thomson Park Drive 
Cranberry Township, PA 16066 
(724) 779-2111 




TABLE OF ELEMENTS, SYMBOLS, ATOMIC NUMBERS, 
AND ATOMIC WEIGHTS 


Element 

Sym¬ 

bol 

At. 

No. 

Atomic 

Weight 

Ap¬ 

prox. 

At. 

Wt. 

Element 

Sym¬ 

bol 

At. 

No. 

Atomic 

Weight 

Ap¬ 

prox. 

At. 

Wt. 

Aluminum 

A1 

13 

26.97 

27 

Mercury 

Hg 

80 

200.61 

200 

Antimony 

Sb 

51 

121.76 

121.5 

Molybdenum 

Mo 

42 

96.0 

— 

Argon 

A 

18 

39.94 

40 

Neodymium 

Nd 

60 

144.27 

— 

Araenic 

As 

33 

74.96 

75 

Neon 

Ne 

10 

20.183 

— 

Barium 

Ba 

56 

137.36 

137 

Nickel 

Ni 

28 

58.69 

58.7 

Beryllium 

Be 

4 

9.02 

— 

Nitrogen 

N 

7 

14.008 

14 

Bismuth 

Bi 

83 

209.00 

209 

Osmium 

Os 

76 

190.8 

— 

Boron 

B 

5 

10.82 

— 

Oxygen 

O 

8 

16.000 

16 

Bromine 

Br 

35 

79.916 

80 

Palladium 

Pd 

46 

106.7 

—- 

Cadmium 

Cd 

48 

112.41 

— 

Phosphorus 

P 

15 

31.027 

31 

Calcium 

Ca 

20 

40.07 

40 

Platinum 

Pt 

78 

195.23 

— 

Carbon 

C 

6 

12.000 

12 

Potas.sium 

K 

19 

39.10 

39 

Cerium 

Ce 

58 

140.13 

— 

Praseodymium 

Pr 

59 

140.92 

— 

Cesium 

Cs 

55 

132.81 

— 

Radium 

Ra 

88 

225.97 

226 

Chlorine 

Cl 

17 

35.457 

35.5 

Radon 

Rn 

86 

222 

— 

Chromium 

Cr 

24 

52.01 

52 

Rhodium 

Rh' 

45 

102.91 

— 

Cobalt 

Co 

27 

58.94 

59 

Rubidium 

Rb 

37 

85.44 

— 

Columbium 

Cb 

41 

93.1 

— 

Ruthenium 

Ru 

44 

101.7 

— 

Copper 

Cu 

29 

63.57 

63.5 

Samarium 

Sm 

62 

150.43 

— 

Dysprosium 

Dy 

66 

162.46 

— 

Scandium 

Sc 

21 

45.10 

— 

Erbium 

Er 

68 

167.7 

— 

Selenium 

Se 

34 

79.2 

— 

Europium 

Eu 

63 

152.0 

— 

Silicon 

Si 

14 

28.06 

28 

Fluorine 

F 

9 

19.00 

19 

Silver 

Ag 

47 

107.880 

108 

Gadolinium 

Gd 

64 

157.26 

— 

Sodium 

Na 

11 

22.997 

23 

Gallium 

Ga 

31 

69.72 

— 

Strontium 

Sr 

38 

87.63 

— 

Germanium 

Ge 

32 

72.60 

— 

Sulfur 

S 

16 

32.064 

32 

Gold 

Au 

79 

197.2 

197 . 

Tantalum 

Ta 

73 

181.5 

— 

Hafnium 

Hf 

72 

178.6 

— 

Tellurium 

Te 

52 

127.5 

— 

Helium 

He 

2 

4.002 

4 

Terbium 

Tb 

65 

159.2 

— 

Holmium 

Ho 

67 

163.5 

— 

Thallium 

T1 

81 

204.39 

— 

Hydrogen 

H 

1 

1.008 

1 

Thorium 

Th 

90 

232.12 

— 

Indium 

In 

49 

114.8 

— 

Thulium 

Tm 

69 

169.4 

— 

Iodine 

I 

53 

126.932 

127 

Tin 

Sn 

50 

118.70 

119 

Iridium v 

Ir 

77 

193.1 

— 

Titanium 

Ti 

22 

47.90 

— 

Iron 

Fe 

26 

55.84 

56 

Tungsten 

W 

74 

184.0 

— 

Krypton 

Kr 

36 

82.9 

— 

Uranium 

U 

92 

238.14 

— 

Lanthanum 

La 

57 

138.90 

— 

Vanadium 

V 

23 

50.96 

— 

Lead 

Pb 

82 

207.22 

207 

Xenon 

Xe 

54 

130.2 

— 

Lithium 

Li 

3 

6.940 

— 

Ytterbium 

Yb 

70 

173.6 

— 

Lutecium 

Lu 

71 

175.0 

— 

Y ttrium 

Y 

39 

88.93 

— 

Magnesium 

Mg 

12 

24.32 

24 

Zinc 

Zn 

30 

65.38 

65 

Manganese 

Mn 

25 

54.93 

55 

Zirconium 

Zr 

40 

91.22 

































